Chemistry NCEA L 2 2 4 Bonding Structure

Chemistry NCEA L 2 2. 4 Bonding, Structure and Energy 1

Achievement Criteria Bonding, structure and energy changes are limited to: Lewis structures, shape and polarity of simple molecules. Simple molecules have no more than four electron pairs about any atom (including multiple-bonded species) Intermolecular forces (the distinction between the different types of intermolecular forces is not required) Ionic, covalent and metallic bonding Molecular, ionic, metallic and covalent network substances Properties are limited to hardness (including malleability and ductility), electrical conductivity, melting and boiling points and solubility. Exothermic and endothermic reactions including energy changes associated with the making and breaking of chemical bonds, differing amounts of substances and changes of state. These may involve calculations. 2

Introduction Chemistry is the study of matter and energy and the interaction between them. The elements are the building blocks of all types of matter in the universe. Each element is made up of only one type of atom, each with its specific number of protons known as its atomic number. 3

Introduction A large amount of energy is required to break an atom down into smaller particles. The elements occur in widely varying quantities on earth. The ten most abundant elements make up 98% of the mass of earth. Many elements occur only in traces, and a few elements are synthetic and highly unstable. 4

Atomic Theory Atoms are the building blocks of elements. Scientists and philosophers have guessed that all matter is made up of building blocks for a very long time. Discovery of the actual structure of the atom has only been in relatively recent times however.

Atomic Theory Democritus – Greek philosopher 400 B. C. The word atom comes from the Greek word atomos meaning indivisible. Democritus reasoned that you could continue to half a piece of matter only to a particular point – the smallest particle that could no longer be divided was known as an atom. The structure of the atom was unknown at this time. Extra for experts

Atomic Theory Extra for experts Boyle – English physicist 1661 Robert Boyle investigated the elements, gold and silver, and found they could not be changed into earth, air, fire or water, that previously Aristotle (Greece 350 BC) claimed all matter was made of. He concluded that gold and silver were true elements and created a definition for an element: An element is any substance that cannot be broken down into a simpler substance. SJ Gaze

Atomic Theory Extra for experts Lavoisier – French Chemist 1743 - 1794 Antoine Lavoisier used a weighing balance to investigate mass relationships in chemical reactions. The law of Conservation of Mass is based on his findings. The total quantity of matter before and after a chemical reaction is the same.

Atomic Theory Extra for experts Proust – French Chemist 1799 Joseph Proust investigated individual elements that made up compounds. The law of Constant Composition is based on his findings. All pure samples of a given compound contain the same elements combined in the same proportion by mass.

Extra for experts Atomic Theory Dalton – English Chemist 1808 John Dalton lived most of his life as a teacher and public lecturer, but his curiosity of the world around him led him into research and experimentation in many areas of Science. One of his most valuable areas of study was in Chemistry concerning the atom and elements. Dalton devised his own notation for elements known at the time.

Atomic Theory Dalton – English Chemist 1808 John Dalton combined previous ideas about atoms in a published a book, called A New System of Chemical Philosophy, adding his own ideas. A summary of the main ideas from his Atomic Theory are: All matter is made up of extremely small particles called atoms. All atoms of one element are identical. Atoms of one element differ in size and mass from atoms of another element. Atoms are indestructible and merely rearrange themselves during chemical reactions. Atoms of different elements combine in simple whole number ratios to form chemical compounds.

Dalton’s Atomic Theory 1. All Matter is made up of atoms 2. Matter can neither be created nor destroyed 3. Atoms of a particular element are alike 4. Atoms of different elements are different 5. A chemical change (reaction) involves the union or separation of individual atoms

Atomic Theory Dalton – English Chemist 1808 Dalton also investigated compounds and came up with the following ideas: Compounds are composed of molecules Molecules are composed of atoms in definite proportions

Atomic Theory Avogadro – Chemist 1811 Amedeo Avogadro also demonstrated that some elements, such as hydrogen and oxygen, existed in their natural state as more than one atom bonded together. He called those particles molecules. A molecule is the smallest particle of a particular substance which can exist in a free state.

Atomic Theory Berzelius – Swedish professor 1813 JÖns Berzelius developed a system of chemical symbols to represent elements. Dalton had devised little pictures but these were difficult to draw. He used letters of the alphabet for the latin name of each element, and a further letter if there was more than one element that started with the same name. Berzelius’s system is now used universally by scientists. The first letter is always a capital letter The second (if there is a second) is always a lower case letter

Atomic Theory Elements consist of only one type of atom. Each element can be represented by a chemical symbol. Most symbols are one or two letters, formed from the name of the element, e. g. Hydrogen H, or Helium He. The first letter of the symbol is always a capital letter. Any other letters are lower case. e. g. Helium is He not HE If the symbols are not based on a elements English name then it is most likely to be based on it’s Latin name, the original language of Science.

Atomic Theory Extra for experts JJ Thompson – English Physicist 1897 Thompson used cathode ray tubes to discover the presence of negatively charged particles 1/2000 the mass of atoms. These particles, called electrons, were found in all atoms. He developed the ‘plum pudding’ model to explain that the negative electrons must be balanced by positive particles to produce a neutral atom.

Atomic Theory Rutherford – New Zealand Physicist 1911 Ernest Rutherford discovered the true structure of atoms by using alpha particles (helium nuclei) fired at gold foil. A large number of these particles travelled straight through and he surmised: Almost all of the space occupied by an atom must be empty From the small number of particles that bounced directly back he surmised: The centre of an atom must be very dense and small From the way the positive alpha particles were deflected is surmised: The centre of an atom must contain a region with a very great positive charge

Atomic Theory Extra for experts Rutherford – New Zealand Physicist 1911 Results from Rutherford’s gold foil experiment could not be explained by the ‘plum pudding’ model of Thompson so Instead, in 1911, Rutherford proposed a new model of the atom in which all of the positive charge is condensed into a tiny, massive nucleus about ten thousand times smaller than the entire atom. Rutherford explained the much lighter electrons circulated outside the nucleus. This was a revolution in the ideas of atoms as Rutherford’s model implied that matter consisted almost entirely of empty space.

Atomic Theory Ernest Rutherford named the positive particles protons found in the dense centre nucleus. Because this accounted for only around half the mass in most atoms he deduced there must be a similar amount of neutral particles - he named these neutrons. Electron -ve Nucleus Neutron N NP P N N P P P NP N Proton +ve Shell 1 Shell 2

Atom Charges Protons are positively charged; electrons are negatively charged; neutrons have o electrical charge. Atoms have no overall charge because the number of protons = number of electrons. All matter is made up of atoms. Atoms consist of protons, neutrons and electrons. The charges of protons and electrons are equal and opposite.

Atom Summary Subatomic particle symbol Mass compared to a proton charge location Proton p 1 +1 In the nucleus Neutron n 1 0 In the nucleus Electron e 1/1840 -1 Moving outside the nucleus

Atomic Theory Moseley – English physicist 1914 Henry Moseley worked alongside Rutherford in investigating atom structure. He discovered some elements had different atomic masses even though they had the same number of protons. He realised elements could be found with different numbers of neutrons and called these varieties isotopes. Atoms with the same number of protons but different numbers of neutrons are called isotopes.

Isotopes of elements occur when atoms have the same atomic number (Z) but different numbers of neutrons in the nucleus. The numbers of neutrons in an atom does not affect the way an element behaves chemically, but it does affect the way it behaves physically. Isotopes found in nature are generally stable, however radioactive isotopes do exist such as 238 Uranium 24

Atomic and Mass number The atomic number is unique for each element. A neutral atom has the same number of electrons as protons. The periodic table is arranged in order of an elements atomic number. The mass number is the total number of protons and neutrons together.

Calculating protons, neutrons and electrons Number of protons: For an atom or ion = atomic number Number of electrons: For an atom = atomic number For a negative ion = atomic number + charge (- =1, -2 =2 etc) For a positive ion = atomic number – charge (+ =1, +2 = 2 etc) Number of neutrons: For an atom or ion = mass number - atomic number atom or ion number of protons number of electrons number of neutrons Mg 12 12 12 Mg 2+ 12 10 12 F 9 9 10 F- 9 10 10

Atomic Theory Bohr – Danish Physicist 1885 - 1962 Niels Bohr was a physicist who also worked alongside Rutherford. He devised a new theory to explain how electrons continuously orbit the nucleus and not gradually lose energy and spiral into it. The main points of his theory were: Electrons only occupy certain orbits of fixed energy Electrons that remain in an orbit don’t emit or absorb energy Electrons can move between orbits. By absorbing energy an electron can move to an orbit further away at a higher energy level. By emitting energy (in the form of light) an electron can move into a lower energy level that is closer to the nucleus. The electron orbits are called energy levels or shells. SJ Gaze

Energy Levels in Atoms Electrons move further away from the nucleus as they gain in energy. Energy can be provided in the form of heat or light. 28

Quantum physics and energy levels Extra for experts Why are electrons in shells or energy levels? It all has to do with quantum physics, where particles, in this case electrons can only contain particular packages of energy. The more packages of energy an electron has, the further out shell it sits in from the nucleus. There can be no half packages or quarter packages; only whole packages, thus the word quantum meaning quantity. Classical physics A turtle sitting on a ramp can have any height above the ground- and so, any energy level. Quantum physics A turtle sitting on a staircase can only sit at certain heights, therefore has only certain energy levels.

The electrons in an atom are arranged in a series of energy levels. Electrons move or ‘orbit’ around the nucleus in energy levels or shells. The energy levels further away from the nucleus are able to fit more electrons. The first energy level is filled first, followed by the second and so on until all the electrons (the same number of protons in an atom) have been used. Maximum numbers of electrons in each energy level are: >2 in the first EL (nearest the nucleus) >8 in the second EL >8 in the third EL (before the fourth shell starts to fill) >18 in the fourth EL You need to draw the configurations of the first 20 elements as well as knowing their names and symbols

Valence electrons The electrons in the outermost energy level of an atom are known as valence electrons. These valence electrons are the particles that react with other atoms in a chemical reaction. The number of valence electrons can determine how an atom will react with others.

There is a relationship between the period number and the number of energy levels an atom has. At this time, the maximum number of energy levels(or electron orbitals) for any element is seven. In the periodic table, elements have something in common if they are in the same row. All of the elements in a period have the same number of energy levels. Every element in the top row (the first period) has one energy level for its electrons) All of the elements in the second row (the second period) have two energy levels for their electrons. It goes down the periodic table like that.

Electron configuration A shorthand way of describing the way electrons are arranged in an atom is called the electron configuration. The information for the number of electrons is found by an elements Atomic Number (number of electrons = number of protons in a neutral atom). Each energy level is filled to its maximum capacity, starting with the lowest energy level first (energy level number 1 or M shell). The energy level are separated by a comma. The energy levels are filled until all the electrons are placed. 12 Atomic number The total of the electronic configuration must equal the atomic number in an atom 2, 8, 2 Mg 24 First EL, second EL, third EL 33

Using the Periodic table to write electron configurations Period number gives number of energy levels Last number of group gives electrons in outer energy level. i. e. group 17 - 7 electrons in outer energy level. Step 1. Ca in period (row 4) so has 4 energy levels Ca 2 , 8 , 2 Step 2. Ca in group 2 so has 2 electrons in the outside energy level Step 3. backfill all energy levels with 8 electrons (2 in first) and add commas between each

Extra for experts Atomic Theory Latest theories in atomic structure. Atoms consist of a nucleus and surrounding electron cloud. The electrons are obeying quantum mechanics; their The mechanics of exact position and speed can’t be measured but we can only give areas of probability in which the electrons can be found. Area where the electrons are most likely to be found orbiting around the nucleus

Periodic table Mendeleev – Russian professor of Chemistry 1834 - 1907 Dimitry Mendeleev was a Chemist who created a periodic table based on elements relative atomic mass and placed the elements in groups based on the elements similar properties. Not all of the elements had been discovered at the time he created the table so he left gaps that has subsequently been filled. Groups 3 to 12 were added after Mendeleev’s table – these are called the transition metals Group 18 – the noble gases, were not discovered at that time and were also added after.

Periodic Table The columns (downwards) of a periodic table are called groups. The rows (across) of a periodic table are called periods. Elements in the same group all have the same number of electrons in their outer (or valence) shells. Elements in the same period all have the same number of shells of electrons in their atoms 37

Group 1 elements These elements are called the Alkali Metals. They are all very reactive with air and, especially so, water. The further down the group the more reactive they are. Hydrogen is not included in this as it does not share similar properties with the rest of the elements. 38

Group 2 elements These elements are called the Alkali Earth Metals. They all react with air, but are less reactive than group 1. 39

Group 17 elements These elements are called the Halogens. They are highly reactive, with reactivity decreasing down the table. They all contain 7 valence electrons and readily accept an electron from other atoms. 40

Group 18 elements These elements are called the Inert Gases (sometimes called Noble gases). They all have full valence shells and are very unreactive. This property makes them useful in many practical applications. 41

ions Atoms with filled outer energy levels are the most stable. An atom will gain or lose electrons in order to have a filled valence shell. Ions have different chemical properties and a physical appearance to the elements they originated from Cation Sodium (Na) 11+ Sodium now becomes the sodium ion Na+ Anion Chlorine (Cl) 17+ Chlorine now becomes the chlorine ion Cl-

ions Ions are charged particles. Ions form when atoms gain or lose electrons. An ion is an atom or group of atoms which has gained or lost electrons. Elements are most stable when the outer energy level (valence shell) is full. Elements can lose or gain electrons when they react with other chemicals to form ions and achieve stability.

ions Atoms that lose electrons form positively charged ions, or cations. Atoms that gain electrons form negatively charged ions, or anions. Cation (Cat) Anion (an Iron) + Metals lose electrons to form Cations. They have 1 -3 electrons in their outside shell Non-Metals gain electrons to form Anions. They have 7 -8 electrons in their outside shell.

Ion Chart - Cations 1+ sodium potassium silver ammonium Na+ K+ Ag+ NH 4+ 2+ magnesium iron (II) ferrous copper (II) cupric Mg 2+ Fe 2+ Cu 2+ zinc Zn 2+ barium Ba 2+ Copper (I) cuprous Cu+ Hydrogen H+ lead Pb 2+ Lithium Li+ tin Sn 2+ 3+ aluminium iron (III) ferric Chromium Al 3+ Fe 3+ Cr 3+

Ion Chart - Anions 1 - 2 - chloride Cl- carbonate CO 32 - permanganate Mn. O 4 - oxide O 2 - thiocyanate SCN- sulfide S 2 - iodide I- sulfate SO 42 - hydroxide OH- sulfite SO 32 - hydrogen carbonate HCO 3 - thiosulfate S 2 O 32 - hydrogen sulfide HSO 3 - chromate Cr. O 42 - fluoride F- dichromate Cr 2 O 72 - bromide Br- nitrate NO 3 - hypochlorite OCl- 3 phosphate 46 PO 4 -3

electron configurations of ions – Cations (metals) The Ca atom has 20 protons and 20 electrons so has no charge. It is neutral. The Ca 2+ ion has 20 protons and 18 electrons so has a 2+ charge.

electron configurations of ions – Anions (non-metals) The Cl atom has 17 protons and 17 electrons so has no charge. It is neutral. The Cl- ion has 17 protons and 18 electrons so has a 1 - charge.

Compounds form from two or more different elements bonded together.

Compounds The compounds are often more stable than the elements they originated from and may release this extra energy in the form of heat and/or light when bonding together. There are two main types of bonding holding atoms together in a compound; Ionic and Covalent. 50

Ionic Bonding is where one atom completely takes valence electrons from another to form ions and the resulting negative and positive ions hold together with electrostatic attraction. This type of bonding occurs when a metal and non-metal react and there is a transfer of electrons to form ions. The ions then combine in a set ratio to form a neutral compound with negative and positive charges balanced out.

Ionic compounds are the product of chemical reactions between metal and non-metal ions Some compounds are ionic compounds, since they are made up of cations and anions. Compounds are neutral substances. For ionic compounds, the charges of the positive ions are balanced by the charges of the negative ions. The Anion (F) takes the electrons off the Cation (Li) so their outer energy levels have a stable 8 electrons each. Anions and Cations have a strong electrostatic attraction for each other so they bond together as a compound.

Covalent Bonding is where electrons are shared between neighbouring atoms. This often occurs when two or more non-metals react. No ions are formed and there is no transfer of electrons. The compound formed is neutral with no charge. The valance electrons (electrons in outside energy level) are involved in bonding. These electrons orbit in pairs. The negative charge of the electron pair will attract the positively positive nucleus of other atoms, and this holds the atoms together in a molecule.

Covalent Bonding Extra for experts All covalent bonds are strong. That is it requires a large amount of energy to ‘break’ the bond. However, some covalent bonds are stronger than others. The greater the overlap of valence orbitals (the area the valence electrons orbit the nucleus) the stronger the bond.

Covalent Bonding The electron-pair must lie between the nuclei for the attraction to outweigh the repulsion of the two nuclei. This ‘sharing’ of electrons between atoms creates a covalent bond – giving both atoms the stability of a full outer shell. Covalent bonds are normally formed between pairs of non-metallic atoms. Some covalent bonds involve only one pair of electrons and are known as single bonds. Other covalent bonds can involve two pairs of electrons; double bonds and three pairs of electrons; triple bonds.

Naming compounds Lavoisier – French Chemist 1789 Lavoisier devised a system of naming compounds based on their chemical composition. If the compound is formed between a Metal cation (+ve) and a Non-Metal anion (-ve), then the compound name joins the two names together with the metal name first. Names of the ions need to be remembered. Sodium + hydroxide

Chemical compound formula 1. Write down the ions (with charges) that react to form the compound. Cation comes before Anion. Al 3+ O 2 - 2. Cross and drop the charge numbers. 3. Place brackets around a compound ion. Al 2 O 3 4. If the numbers are both the same remove. 5. If any of the numbers are a 1 they are removed 6. Remove any brackets if not followed by a number H+ SO 4 -2 H 2(SO 4)1 H 2 SO 4

Drawing chemical compounds G Lewis – American Chemist 1916 G Lewis devised a system of drawing covalent molecules showing arrangement of atoms and valence electrons – both those involved in bonding and those that are not (called lone pairs). Electrons in inner shells are not involved in bonding. These diagrams are called Lewis diagrams. The Lewis diagram is drawn so that each atom has eight electrons associated with it (except for hydrogen which has two). This is the octet rule. Lewis diagram of H 2 O (water) Hydrogen electron Bonded pair H O x Oxygen electron x H Lone pair

Lewis Diagrams 1. Calculate valence electrons of all atoms. If the molecule is an ion then subtract the charge from the total electrons and place the charge outside of square brackets of the Lewis diagram at the end. Example carbon dioxide. C=4 CO 2 O=6 16 2. Write down number of pairs of electrons. 16 / 2 = 8 pairs 3. Place atom with least filled valence shell in the centre with the other atoms arranged around the out side (periphery) O C O

Lewis Diagrams 4. Bond all atoms together (either x or O C O = one pair of electrons) 8 pairs – 2 pairs = 6 pairs remaining 5. Place remaining e- pairs around the periphery atoms so each has 4 pairs (including bond pair) around it. xx x x. O xx C xx 6 pairs – 6 pairs = 0 pairs remaining O xx xx 6. If there any remaining pairs place them around the outside of the central atom. 7. Rearrange lone pairs (pairs not bonded) into bonded pairs if the central atom does not have 4 pairs around it.

Rule of orbitals – exceptions to the rule If there are extra Lone Pairs of electrons left after all of the periphery atoms are filled in accordance with the octet rule then they are placed around the central atom(s) according to the Rule of Orbitals. The Rule of Orbitals: the total number of lone pairs and bond pairs (LP+BP) associated with an atom cannot exceed the number of Valence Shell Orbitals (VSO = n 2, where n is the row of the Periodic Table in which that atom resides). n = 1 (H): maximum VSE pairs (LP+BP) = VSO = 1; n = 2 (B, C, N, O, F): maximum VSE pairs (LP+BP) = VSO = 4 ("octet rule") n = 3 ((Al, Si, P, S, Cl): maximum VSE pairs (LP+BP) = VSO = 9 etc. Boron and Beryllium often are found with only 3 lone + bonded pairs around them F B F F

Lewis Diagrams § § § Extra for experts The number of covalent bonds an atom forms is called its valence. Some atoms have fixed valence. E. g: H = 1, C = 4, F = 1. (most halogens = 1) Some atoms have variable valence. For example: O = 2 (sometimes 3), B, N = 3 (sometimes 4). an atom bonded to only one other atom is peripheral (monovalent atoms such as H and F are always peripheral). an atom bonded to two or more other atoms is central. Often, the formula is written to indicate connectivity. For example: HCN = H bonded to C, C bonded to N, H and N are not bonded.

Predicting molecular shapes Sidgwick and Powell – 1940 Sidgewick and Powell devised a theory to predict the shapes molecule formed. It is based on the following ideas: Each electron pair is a region of negative charge Negative charges repel each other Electron pairs will be spaced as far apart as possible around a central atom. This theory is called the Valance Electron Pair repulsion theory. Three pairs Four pairs

Predicting molecular shapes The shapes of molecules are determined by the way the regions of negative charge arranged around the central atom in the molecule. A region may consist of one lone pair of electrons or one bonded pair or two bonded pairs or three bonded pairs. All of these electron arrangements occupy the same region of space

Shapes – two clouds Since regions of electrons are negatively-charged, they repel each other as far apart as possible. Two clouds arrange themselves on opposite sides of the central atom. The bond angle will be 180°. The shape name is linear. Bonded pair of electrons (one cloud) atom 180° 65 Bond angle

Shapes – three clouds (0 lone pairs) Three regions of negative charge will cause a bond angle of 120° as they repel each other. All the atoms still lie on a flat plane (like a sheet of paper). The shape is trigonal planar. (or triangular planar) Periphery atom 120° Central atom

Shapes – two clouds (1 lone pair) When one of the clouds of electrons is a lone pair it will have a slightly greater push to the bonded pairs. This is because the lone pair are only orbiting around one positive nucleus and their negative charge is less ‘neutralised’ than if they had another nucleus to orbit around. The regions of negative charge repel to a trigonal planar shape. The bond angle between the remaining pairs is approximately 120°. The final shape formed by the atoms is called bent. Lone pair Bonded pair ~120°

Shapes – four clouds (0 lone pairs) When four regions of negative charge around a central atom they repel each other into a 3 -dimensional shape. The bond angle is now 109. 5°. This is because it is a sphere divided into 4 rather than a circle. This shape is tetrahedral. 109. 5°

Shapes – four clouds (1 lone pair) The four regions of negative charge still occupy a 3 -dimensional tetrahedral shape. (The lone pair, however, exerts a stronger repulsion to the remaining bonded pairs). The bond angle is 109. 5°. The final shape the bonded atoms form is a trigonal pyramid (or a triangular pyramid) Lone pair 109. 5°

Shapes – four clouds (2 lone pairs) The 4 regions of negative charge repel each other to a (warped) tetrahedral shape. But The two lone pairs create a much stronger repulsion than one lone pair and the bond angle between the remaining bonded pairs is smaller again at approximately 109. 5° (compared to 120 ° of the bent shape with only 3 regions of negative charge). The final shape the bonded atoms form is called Bent. Lone pair 109. 5°

Shape Summary Electron clouds (lone pairs or bond groups) 2 Linear CO 2 180° 3 4 no lone pairs 1 lone pair no lone pair 1 lone pair 2 lone pair triangular planar Bent Tetrahedral triangular pyramid CH 4 109. 5° NH 3 109. 5° BF 3 120° SO 2 ~120° bent H 2 O 109. 5°

Drawing Shapes 1. Draw molecule – Lewis diagram first 2. Calculate number ofregions of negative charge around central atom • • Single, double or triple electron bonds occupy 1 cloud Note: make sure the question asks you to draw a shape and not a lewis diagram receding Same plane A lone pair of electrons occupy 1 cloud 3. Calculate how many atoms are joined to the central atom 4. Name / draw shape approaching

Discussing shapes questions – NCEA example Explain why the shape of the CO 2 molecule is linear but the shape of H 2 O is bent? 1. The C (central atom) of CO 2 has 2 regions of negative charge around it in the form of double bonds connected to a O atom. (draw lewis diagram) 2. Each of the regions of negative charge repel each other the furthest away from each other in 3 dimensional space into a linear. 3. There are no lone pairs so the final CO 2 molecule therefore also forms a linear shape 1. The O molecule (central atom) of H 2 O has 4 regions of negative charge around it in the form of two single bonds connected to a H atom and two lone pairs. (draw lewis diagram) 2. Each of the regions of negative charge repel each other the furthest away from each other in 3 dimensional space and form a tetrahedral shape. 3. However with only 2 of the regions bonded to atoms the final shape the H 2 O molecule forms is a bent shape

Electronegativity is the attraction that an atom has towards electrons from another atom. The greater the electronegativity the stronger the pull it has towards other electrons. Trends in the periodic table § The larger the nucleus (with the positive protons) the stronger the electronegativity, this means it increases from left to right. §The further the valence electrons are from the nucleus the lesser the electronegativity, therefore the electronegativity decreases down a group.

Electronegativity We use a pauling scale to determine electronegativity. The scale starts close to 0 – with minimal electronegativity and goes up to 4 with the highest electronegativity. Most of the Inert gases do not have a value because of there non reactivity with other atoms.

Extra for experts Ionic – covalent bond continuum Bond types between atoms can depend on the electronegativity of the atoms. Rather than discrete categories, molecules fall along a continuum Covalent 0— 0. 4 Polar Covalent 0. 4— 1. 6 Ionic >1. 6

Polarity If two identical atoms are bonded together then they have exactly the same amount of attraction to the shared electrons in the bonded pair. This is because their electronegativity is the same. This becomes a non-polar molecule with non-polar bonds. Example - Iodine molecule I 2 Path of electrons Evenly shared If two different types of atoms are bonded together then they will exert different levels of attraction for the orbiting electrons. That is because they may have different numbers of electron shells and different numbers of protons in their nucleus. This will cause an electronegativity difference. These bonds become polar bonds. Example – hydrochloric acid HCl Slightly positive Electrons orbit less δ+ δ- Slightly negative Electrons orbit more

Polarity δ- δ+ Cl H Polarity may also be shown as an arrow, with a cross, +ve, at the tail. The arrow head is the –ve end.

Polarity If two bonded atoms are the same, the bond is said to be non-polar. i. e. I 2 The whole molecule is also non-polar because there is no electronegativity difference and the valence electrons orbit each atom evenly. If two different atoms are bonded they form a polar bond, as there is an electronegativity difference and the valence electrons spend more time around the atom with the higher electronegativity value (that atom becomes slightly negative ) The atom that the valence electrons spend less time around becomes slightly positive.

Symmetry and Polarity The polarity of a molecule with polar bonds depends upon whether the molecule is symmetrical or not. A symmetrical molecule (one where the centres of peripheral atoms coincide) becomes a non -polar molecule – as the charges balance out An unsymmetrical molecule (one where the centre of peripheral atoms do not coincide) is a polar molecule.

Answering NCEA Polarity Questions Explain why molecules x (CCl 4) and y (NCl 3) are polar and non-polar? Polar molecule 1. molecule (NCl 3) is polar (state which one) 2. (NCl 3) contains polar bonds due to electronegativity difference of N and Cl. 3. over the whole molecule the atoms are not distributed symmetrically in 3 dimensions because its shape is (state which one) and has lone pairs of electrons 4. polar bonds do not cancel each other out and the whole molecule is polar. Non-polar molecule 1. molecule (CCl 4) is non-polar (state which one) 2. (CCl 4) contains polar bonds due to electronegativity difference of C and Cl. Cl attracts more electrons than C because it has a bigger atomic number than C but with the same number of shells 3. over the whole molecule the atoms are distributed symmetrically in 3 dimensions because its shape is (state which one) 4. polar bonds cancel each other out and the whole molecule is non-polar.

Solubility The solubility of a substance is the amount of that substance that will dissolve in a given amount of solvent. Solubility is a quantitative term. Solubility's vary depending on the solvent and the solute. The terms soluble and insoluble are relative. Some substances can be sparingly soluble where only the most minute percentage dissolves. For a solute to dissolve the attraction to the solvent molecules must be stronger than the bonds holding the atoms/molecules of the solute together. GZ Science Resources 2013

Water as a solvent Solute (salt) Solvent (water) Solution (saltwater) A solution is made up of a solvent and a solute. A solvent is a substance such as water that is able to dissolve a solute. The solvent ‘pulls apart’ the bonds that hold the solute together and the solute particles diffuse (spread randomly by hitting into each other) throughout the solvent to create a solution. The solution is a mixture with evenly spread solvent and solute particles. These particles can be physically separated by evaporation.

Solutions form when a solute is dissolved in a solvent When a solid mixes into a liquid and can longer be seen it has dissolved. The liquid is called the solvent and it pulls apart the bonds between the solid particles, called the solute, and they diffuse. A solution is then created when the solvent particles (often water) are mixed up with the broken apart solute particles.

Solubility For a solute to dissolve, the solvent particles must form bonds with the solute particles that are of similar strength, to the bonds between the solute particles. Water, being polar attracts ions because they are charged and so dissolves many ionic substances.

Polar Solvents The water molecule has two polar bonds. Due to the asymmetry of the molecule, their polarities reinforce making the oxygen side of the molecule partially negative (δ-) and the hydrogen side partially positive (δ+). Such molecules are called ‘polar’. The separation of charge in the molecule (δ+ - δ) is called a ‘dipole’. Polarity causes a stream of water molecules to attract to a charged plastic pen.

Dissolving and Polarity Polar substances dissolve polar substances. e. g. Water, being polar attracts the molecules of other polar substances (e. g. HCl) and will dissolve them. Polar substances will not dissolve non-polar substances. e. g. Water, (polar) has a stronger attraction to itself than to non-polar molecules (e. g. cyclohexane) and will not dissolve them. Non-polar substances dissolve non-polar substances. e. g. Non polar solvents (like cyclohexane) attract non-polar solutes (like napthalene) by the same weak Van der Waals forces they attract themselves by and so will dissolve non-polar solutes. Ionic solid dissolving in water

Common Polar and Non-polar molecular substances Polar Non-Polar water methanol acetic acid hydrogen chloride cyclohexane benzene hydrocarbons (e. g. petrol) oxygen hydrogen nitrogen iodine Sample NCEA Questions: Potassium chloride will not dissolve in non-polar solvents, but will dissolve in water. Explain by relating the property to the structure and bonding within the solid. Use the structure and bonding in H 2 O and SO 2 to explain why SO 2 is soluble in H 2 O.

Groups of substances Substances are grouped together according to the type of bonds they have between particles. This year will cover four groups of substances; Molecular, metallic, ionic and covalent network. The physical properties of these groups will be linked to their structure. Molecular solids Ionic solids Non-metals forming molecules Non-metals and metals together forming a ionic compound S 2 sulfur HCl Hydrogen chloride I 2 iodine KI Potassium iodide Cu. SO 4 Na. Cl Copper sulfate Sodium chloride Metallic solids Covalent network solids Elements that are metals Carbon and silicon dioxide Fe iron Si. O 2 Silicon dioxide Al aluminium Cu copper C diamond C graphite

Non-polar Molecular solids non-metal + non-metal Molecules are held together by weak intermolecular forces caused by temporary dipoles induced by electrons randomly spending more time around one nucleus than the other. Within the Molecules, the atoms are held together by strong covalent bonds. Weak intermolecular bond Strong covalent bond

Polar Molecular solids Polar molecules held together by weak inter molecular forces caused by permanent dipoles induced by electrons spending more time around one nucleus in the molecule that has greater electronegativity than the other. The δ –ve end of one molecule is attracted to the δ +ve end of another. δ +ve δ -ve δ +ve 91 δ -ve δ +ve

Weak intermolecular forces of attraction 3 kinds instantaneous dipole – induced dipole in non-polar molecules permanent dipole – permanent dipole in polar molecules hydrogen bonding – strong dipoles for example in water between H and O Note the distinction: Intra-molecular Forces: the strong bonding forces within a molecule. i. e. the covalent bonds holding the molecule together. Inter-molecular Forces: the weak bonding forces between molecules due to the attractions between partial charges. i. e. permanent dipole

Polar Molecular solids - solubility 1. Hydrogen chloride (HCl) is a molecular solid 2. Hydrogen chloride is made up of covalently bonded atoms to form molecules 3. these molecules are held together by weak intermolecular forces 4. these molecules are polar therefore the electrostatic attractions of water molecules (which is stronger than the weak intermolecular forces) have sufficient strength to pull the molecules apart hence hydrogen is soluble Polar molecules H Cl Cl H H Cl δ -ve H δ +ve O H δ +ve

Non-polar Molecular solids - solubility 1. Iodine is a molecular solid 2. Iodine is made up of covalently bonded atoms to form molecules 3. these molecules are held together by weak intermolecular forces 4. Iodine is non-polar therefore the electrostatic charges of the water do not have sufficient strength to overcome the weak intermolecular forces holding the molecules together hence iodine is insoluble + + δ -ve H δ +ve O + + H δ -ve O δ +ve H δ +ve Non- Polar molecules H δ +ve

Molecular solids – Melting point For example: 1. Carbon dioxide is a molecular solid (at low temperatures below -56◦C) 2. Carbon dioxide is made up of covalently bonded atoms to form molecules 3. these molecules are held together by weak intermolecular forces 4. these forces require small amounts of energy to break apart the solid (but not the individual molecules which are held together by strong covalent bonds) therefore carbon dioxide has low melting point O Heat O C O Many molecular solids are only solid at temperatures well below 0◦C and at room temperature they are gases

Molecular solids - hardness For example: 1. sulfur is a molecular solid 2. sulfur is made up of covalently bonded atoms to form molecules 3. these molecules are held together by weak inter molecular forces 4. these forces require small amounts of energy to break apart the solid (but not the individual molecules which are held together by strong covalent bonds) therefore sulfur is easily broken up S S S Force applied S S Weak intermolecular force

Molecular solids - Conductivity In order for a substance to be electrically conductive there must be free moving charged particles For example: 1. Iodine is a molecular solid 2. Iodine is made up of covalently bonded atoms to form molecules these molecules are held together by weak intermolecular forces 3. there are no free moving charges therefore iodine cannot conduct electricity nucleus Valence electrons Covalently shared electrons + + Weak intermolecular bonding Fully occupied valence electrons remain in ‘fixed orbit’ around nucleus and are not available to carry charge. The molecule is neutral

Metallic Solids - structure Metals atoms are arranged as positive ions held in place in ordered layers by strong attractive non-directional bonding, forming a lattice. - this gives metals strength. Metal atoms are held together in a 3–D lattice by metallic bonding in which valence electrons are attracted to the nuclei of neighbouring atoms. The attraction of the metal atoms for the valence electrons is not in any particular direction; therefore metal atoms can move past one another without disrupting the metallic bonding, therefore metal is ductile. The atoms are packed tightly together - this makes metals dense

Metallic Solids - Conductivity Electrons from the outer shells of the metal atoms move freely throughout the lattice. - this makes metals excellent conductors of heat and electricity 99

Metallic Solids - Conductivity Free moving charged particles are required to carry a charge and for a substance to be electrically conductive For example: Always state what type of solid a substance is first. 1. copper is a metallic solid 2. copper is arranged as positive ions held in place in ordered layers by strong attractive non-directional forces, in a sea of de-localised electrons 3. electrons are free moving hence can carry a charge 4. therefore copper can conduct electricity

Metallic Solids - Solubility For example: 1. lead is a metallic solid 2. Lead is arranged as positive ions held in place in ordered layers by strong attractive non -directional forces, in a sea of de-localised electrons 3. these forces require a large amount of energy to break therefore the electrostatic attractions of water molecules do not have sufficient strength to pull the atoms apart 4. therefore lead is insoluble In order for substance to dissolve in water (a polar liquid) the attraction between the particles in a substance must be less than the attraction towards water molecules

Metallic Solids – Malleability and ductility For example: 1. iron is a metallic solid 2. iron is arranged as positive ions held in place in ordered layers - a lattice, by strong attractive non-directional forces, in a sea of de-localised electrons 3. these forces require large amounts of energy to break apart the solid therefore aluminium is not easily broken up 4. However Layers can slide over each other, and as the attractive forces are nondirection the metallic particles remain strongly bonded. – this gives the metallic solids the properties of being malleable (moulded into flat sheets) and ductile (drawn out to thin wires) Layers of ions can slide over each other without breaking- this makes metals hard and also malleable and ductile

Metallic Solids – melting point The strength of the bonds between particles determines the energy required to break them, and therefore the amount of energy to change a solid into a liquid (the melting point) where the bonds are somewhat broken. Metals in general, have very strong bonds which makes them solid at room temperature (Mercury is the exception) For example: 1. Aluminium is a metallic solid 2. Aluminium is arranged as positive ions held in place in ordered layers by strong attractive non-directional forces, in a sea of de-localised electrons 3. These forces require a large amount of energy (high temperature) to break apart the metallic solid therefore the melting point is very high. Three steps to answering structure and physical properties questions. The first is state the name of the solid. The second is describe the structure of the solid. The third is link the structure of the solid to the physical property discussed.

Ionic Solids - structure Metal + Non-Metal An ionic solid is made up of ions held together by strong electrostatic forces (ionic bonding) between +ve (cations) and –ve (anions) ions in a 3 -d lattice.

Ionic Solids - Solubility In order for substance to dissolve in water (a polar liquid) the attraction between the particles in a substance must be less than the attraction towards water molecules For example: 1. Sodium chloride (Na. Cl) is an ionic solid 2. Sodium chloride is made up of ions held together by strong electrostatic attractions between +ve and –ve ions in a lattice 3. the electrostatic attractions of water molecules have sufficient strength to pull the ions apart however 4. therefore the solid will dissolve and is soluble Na. Cl first place in water Na+ and Cl- ions breaking apart The positive hydrogen end of water is attracted to the anions and the negative oxygen end of water is attracted to the cations

Ionic Solids - Conductivity Free moving charged particles are required to carry a charge and for a substance to be electrically conductive For example: 1. Sodium chloride is an ionic solid 2. Sodium chloride is made up of ions held together by strong electrostatic forces between +ve and –ve ions in a 3 -d lattice 3. when solid the ions are not free to move therefore it doesn’t conduct electricity 4. However when melted, or dissolved in solution, the bonds are broken and the ions are free to move therefore sodium chloride can conduct electricity Na+ Cl- Cl- Na+

Distilled water does not conduct a current Positive and negative ions fixed in a solid do not conduct a current In solution, positive and negative ions move and conduct a current

Ionic Solids – Hardness and brittleness For Example: 1. Sodium chloride is an ionic solid 2. Sodium chloride is made up of ions held together by strong electrostatic attractions between +ve and –ve ions in a 3 -d lattice so requires a lot of energy to break the bonds 3. However if sideways force is applied and a sheet of the lattice slides then ions of the same charge may come in contact with each other and repel hence the ionic solid is brittle (and can break into pieces) Na+ Cl- Cl- Na+ force

Ionic Solids – Melting Point For example: 1. sodium chloride is an ionic solid 2. Sodium chloride is made up of ions held together by strong electrostatic attractions between +ve and –ve ions in a 3 -d lattice 3. Because these strong bonds require a large amount of energy to break the ionic solids have a high melting point. Na+ Cl- Cl- Na+

Covalent Network Solids - structure All atoms are held together by strong covalent bonds Diamond is a 3 -dimensional covalent network structure where atoms are held together by strong covalent bonds in all planes diamond graphite Graphite is a covalent network structure that is in 2 dimensional sheets (graphite). Between the layers are free moving electrons from the valance electrons of the carbon atoms. Silicon dioxide (Si. O 2) is a 3 dimensional covalent network structure Silicon dioxide

Covalent Network (3 D) - Conductivity of diamond For example: 1. diamond is a 3 dimensional covalent network structure (diamond) 2. all atoms are held together by strong covalent bonds 3. there is no free moving charged particles 4. therefore diamond cannot conduct electricity

Covalent Network (2 D) - Conductivity of graphite For example: 1. graphite is a covalent network that is in 2 dimensional sheets 2. between the layers are free moving electrons from the valance electrons of the carbon atoms. 3. the free moving electrons can carry a current 4. therefore graphite can conduct electricity

Covalent Network Solids - Solubility For example: 1. Silicon Dioxide is a 3 dimensional (or 2 -dimensional) covalent network structure 2. all atoms are held together by strong covalent bonds 3. these forces require a large amount of energy to break therefore the electrostatic attractions of water molecules do not have sufficient strength to pull the ions apart 4. hence silicon dioxide will not dissolve in water and is insoluble

Covalent Network Solids - Melting Point For example: 1 Diamond is a 3 -dimensional (or 2 -dimensional) covalent network structure 2. all atoms are held together by strong covalent bonds 3. these forces require a large amount of energy to break 4. therefore diamond has a very high melting point.

Covalent Network Solids - Hardness For example: 1. Diamond is a 3 -dimensional (or 2 -dimensional) covalent network structure 2. all atoms are held together by strong covalent bonds 3. these forces require a large amount of energy to break 4. therefore diamond is very hard.

Solids Summary Name of solid substance Type of particle in solid Attractive force broken when solid melts Attractive force between particle – weak or strong (hardness) Relative melting point solubility Electrical conductivity molecular molecules Weak intermolecular weak low Yes if polar No if nonpolar no metal atoms Metallic bonding strong high no yes ionic ions Electrostatic strong high yes Only if molten or in solution high no no Ionic bonding covalent atoms Covalent bonding strong

Property Type of Solid Molecular Solubility in Water Electrical Conductivity Melting Point Hardness Covalent network Ionic Metallic

Energy Changes Enthalpy and Enthalpy Change ∆H Enthalpy (or Heat Content) is the energy in a substance due to kinetic energy of particles and potential energy in chemical bonds Enthalpy change ∆H is the difference in enthalpy of products HP and reactants HR ∆H = HP - HR The unit for Enthalpy is kilo. Joules (k. J)

Enthalpy Change HP (products) and HR (reactants) cannot be measured. We can measure Enthalpy change ( ∆H ) by measuring energy; Released to surroundings (Exothermic Reactions) Absorbed from surroundings (Endothermic Reactions) 119

Exothermic Reactions These are reactions where heat energy is released into the surroundings. Surroundings gain heat energy. (increase in temperature ) Products will have less energy than reactants. ∆H is NEGATIVE (-) Endothermic Reactions These are reactions where heat energy is absorbed from the surroundings. Surroundings lose heat energy. (Decrease in temperature) Products will have more energy than reactants. ∆H is POSITIVE (+)

Exothermic reactions Any combustion reaction is exothermic. The bonds holding the atoms of fuel molecules together (usually consisting of carbon and hydrogen atoms) release a lot of energy in the form of light and heat when they are broken. The total energy holding the bonds together in the products are less than the total energy in the reactions and the difference is released. GZ Science Resources 2013 121

Endothermic reactions Melting ice is an example of an endothermic reaction. The solid ice (water) atoms that are in a fixed pattern are barely moving and need to absorb energy in order to move faster and break the bonds to form water in a liquid state. GZ Science Resources 2013 122

Enthalpy Diagrams Endothermic Reactions e. g. Reacting methane with steam at high pressure and temp. Energy is absorbed Exothermic Reactions e. g. Burning of methane in air. Energy is released

Enthalpy Change An exothermic reaction will release energy and the products will be at a lower enthalpy level than the reactants. The reaction system will feel hot to the touch as the energy is released as heat energy. An endothermic reaction will absorb energy and the products will be at a higher enthalpy than the reactants. The reaction system will feel cool to the touch as heat energy is taken from the surroundings, including your skin, and used to break bonds in the molecules.

Energy Diagrams Endothermic Reaction e. g. Reacting methane with steam at high pressure and temp. Energy is absorbed

Energy Diagrams Exothermic Reaction e. g. concentrated Hydrochloric acid reacting with zinc metal HR HP

Breaking Bonds - endothermic Bonds holding atoms and molecules together require the input of energy in order to break them apart therefore breaking of bonds is an endothermic reaction. The input of energy (usually light or heat energy) cause the atoms and molecules to move faster and ‘pull away’ from each other. Each type of bond has its own specific amount of energy, called bond energy measured in k. J, required to break its bond.

Forming Bonds - exothermic Bonds forming between atoms and molecules release energy therefore bond forming is an exothermic reaction. Bonds are formed to form a stable molecule. If more energy is required to break the bonds of the reactants than released when the bonds of the products are form then the overall reaction is endothermic. If less energy is required to break the bonds than is released when the bonds of the products are formed then the overall reaction is exothermic.

Enthalpy in Dissolving If more energy is released when water bonds to the solute than it takes to separate the solute, the dissolving is exothermic and the temperature increases. An example is adding a strong acid (such as sulfuric acid) or base (such as sodium hydroxide)

Standard conditions Measurements depend on conditions When measuring a enthalpy change you will get different values under different conditions. For example, the enthalpy change of a particular reaction will be different at different temperatures, different pressures or different concentrations of reactants. The values for enthalpy are given for standard conditions, indicated by the superscript θ Standard conditions include: Temperature of 25°C Atmospheric pressure conditions of 1 ATM Concentration of 1 mol per Litre

Thermochemical Equations Exothermic CH 4 + 2 O 2 Endothermic CO 2 + 2 H 2 O CH 4 + H 2 O ∆H = -888 k. Jmol-1 CO + 3 H 2 ∆H = 206 k. Jmol-1 This thermochemical equation reads; 888 k. J of heat is released when 1 mole of CH 4 reacts with 2 moles of O 2 to produce 1 mole of CO 2 and 2 moles of H 2 O Use thermochemical equations to find This thermochemical reaction reads; 206 k. J of heat is absorbed when 1 mole of CH 4 reacts with 1 mole of H 2 O to produce 1 mole of CO and 3 moles of H 2. ∆H, n and m. n = m/M n = moles (6. 02 x 1023 particles) m = mass (grams) M = Molar Mass (gmol-1 )

Thermochemical Equation Example CH 4 + 2 O 2 CO 2 + 2 H 2 O ∆r. H = -888 k. Jmol-1 Use the equation above to find heat released if 2. 5 moles of CH 4 burns. 1 mole of CH 4 releases 888 k. J 2. 5 moles CH 4 releases x k. J x = 2. 5 x 888 = 2220 k. J An equation is a mole ratio – the number in front of each substance tells you how many moles of that there is to any other substance. For example there is 1 mole of CH 4 to every 2 moles of O 2 The enthalpy of the equation shows you the amount of energy per unit of substance. 888 = 1 CH 4 888 = 2 O 2 (444 = 1 O 2)

Thermochemical Equation Example 2 CH 4 + 2 O 2 CO 2 + 2 H 2 O ∆r. H = -888 k. Jmol-1 Calculate the amount (in moles) of H 2 O produced when the reaction above releases 10, 000 k. J. An alternative method is to find out how much energy is released per mole first 2 moles H 2 O = 888 k. J Therefore 1 mole H 2 O = 444 k. J Amount of mols in equation 10, 000 k. J/444 k. J = 22. 5 Amount energy per unit of substance Total energy released So 22. 5 moles of water are produced at 444 k. J to reach 10, 000 k. J (22. 5 x 444 = 10, 000)

Thermochemical Equation Example 3 CH 4 + H 2 O CO + 3 H 2 ∆r. H = 206 k. Jmol-1 M(C) = 12 gmol-1 M(O) = 16 gmol-1 Calculate the energy required to produce 1 kg of CO gas from the reaction above Step one moles of CO produced M = 1000 g M(CO) = 28 gmol-1 n = m/M n = 1000/28 = 35. 7 moles 1 kg = 1000 g. Must be converted to grams first If Molar mass is not given then use the periodic table The units are k. J not k. Jmol-1 as it is total amount not amount per mole. Step two 1 mole CO produced requires 206 k. J (as per the equation above) 35. 7 mols CO produced so…. . enthalpy = 35. 7 x 206 = 7354 k. J

Bond Enthalpy The high values for bond enthalpy explains why some substances are very resistant to chemical attack and form very stable molecules In a polyatomic (more than one atom) molecule, the bond strength between a given pair of atoms can vary slightly from one compound to another. The value given for bond enthalpy is the average of all these variations. A multiple bond (double/triple) is always stronger than a single bond because more electrons bind the multiple bonded atoms together. The table below shows some common average bond enthalpies. C= C

Bond Energies Bonds Broken – Endothermic Bonds formed – Exothermic ∆r. H°= ∑ (energy of bonds broken) - ∑(energy of bonds formed) Note: Bond energies calculated for gases. Convert using ∆vap. H° or ∆sub. H° if in solid or liquid state. CO(g) + H 2 O(g) → H 2(g) + CO 2(g) Bonds Broken CΞO H-O x 2 995 k. J 2(463)k. J 1921 k. J ∆r. H°= -1. 0 Kjmol-1 Bonds formed C=O x 2 H-H 2(743)k. J 436 k. J 1922 k. J ∆r. H°= 1921 k. Jmol-1 – 1922 k. Jmol-1 The equation can also be arranged ∆r. H°= -1. 0 k. Jmol-1 to calculate unknown bond energy

Bond Energies Bonds Broken – Endothermic Bonds formed – Exothermic ∆r. H°= ∑ (energy of bonds broken) - ∑(energy of bonds formed) Reactants: Draw lewis diagrams to calculate the number and type of bond Multiply the bond energy given by the number of bonds Products: Draw lewis diagrams to calculate the number and type of bond Total the bond energy for product molecules Total the bond energy for reactant molecules bonds broken (reactants) minus bonds formed (product) = total enthalpy

Using Bond Enthalpies to calculate ∆r. H° Bond enthalpy is the change in enthalpy when the covalent bond, in a gaseous molecule, is broken. It is always a positive value because bond breaking always requires an input of energy. Making bonds releases energy so generally speaking the more bonds a substance can form the more stable it will be. The strength of a covalent bond depends on the electrostatic attraction between the positive nuclei and the shared electron pair. The larger the atomic radius of an atom (which increases down a group) the further the shared electron pair from the positive nucleus – which creates decreasing electrostatic attraction. Therefore the weaker the covalent bond and the lower the value of bond enthalpy The stronger a covalent bond, the higher the value of the bond enthalpy. The units are k. J mol-1.

Calculating ∆r. H° given the standard heats of formation of reactants and products. The standard enthalpy of any reaction can be obtained by subtraction of the standard enthalpies of formation of reactants from those of the products. r. Ho = n f. Hoproducts - n f. Horeactantss where n is the stoichiometric coefficient of each substance in the reaction equation. Example Using the standard heats of formation of CO 2(g), H 2 O(l), and C 6 H 12 O 6(s), calculate the standard enthalpy of combustion of glucose. f. Ho(C 6 H 12 O 6, s) = -1268 k. J mol-1 f. Ho(CO 2, g) = -394 k. J mol-1 f. Ho(H 2 O, l) = -286 k. J mol-1 f. Ho(O 2, g) = 0 k. J mol-1 Note - Start by writing an equation for the combustion of 1 mole of glucose. C 6 H 12 O 6(s) + 6 O 2(g) 6 CO 2(g) + 6 H 2 O(l) r. Ho = n f. Hoproducts - n f. Horeactants r. Ho = ( 6 x -394 + 6 x -286) - (1 x -1268 + 6 x 0) = - 2812 k. J mol-1

Calculating enthalpy using temperature change Extra for experts ΔH can be calculated if you can measure the temperature change (°C) of a particular mass (g) ΔH = m x c x ΔT ΔH = mass (g) x ΔT (°C) x specific heat capacity (c) (c = 4. 18 J g-1 °C) 1. Calculate the mass of reactants – record in grams (one ml liquid = one gram) 2. Measure the temperature of the reactants and the temperature of the products and calculate ΔT (°C) 3. Calculate ΔH using formula above

Enthalpy of Reaction using calorimetry Extra for experts To measure enthalpy changes, the reaction it is carried out in an insulated container (such as a polystyrene cup) and the temperature change (in °C) is measured. Using this temperature change, ΔT, and the value of the specific heat capacity, c, the amount of energy transferred to the mass m of substance (usually water) can be calculated using the expression ∆H = m c ΔT The specific heat capacity of the water is 4. 18 J °C-1 g-1. Every 1 m. L of water can be taken as 1 g due to its density 1. Calculate the mass of reactants – record in grams (one ml liquid = one gram) 2. Measure the temperature of the reactants and the temperature of the products and calculate ΔT (°C) 3. Calculate ΔH using formula above

Specific Heat Capacity (c) The heat required to produce a 1 o. C rise in 1 kg of a substance. Extra for experts e. g. c(H 2 O) = 4. 18 k. J o. C-1 kg-1 The energy change when a body of mass m experiences a temperature rise of ΔT is given by: Q = m c ΔT energy change = mass x specific heat capacity x change in temperature Example: Calculate the energy change when a 9 kg mass of water increases its temperature by 20 o. C. Q = m c ΔT = 9 kg x 4. 18 k. J o. C-1 kg-1 x 20 o. C = 752 k. J Note: Watch units g or kg, J or k. J and be ready to convert if required. 142

Enthalpy Changes Exothermic - energy released ∆ H negative Endothermic - energy absorbed ∆ H positive Enthalpy of reaction (∆r. H) for any given reaction Standard conditions (∆r. H°) 1 atmosphere of pressure, 25°C Enthalpy of fusion (∆fus. H°) 1 mol solid to liquid state Enthalpy of vapourisation (∆vap. H°) 1 mol liquid to gas state Enthalpy of sublimination (∆sub. H°) 1 mol solid to gas state solid ∆sub. H° ∆fus. H° gas liquid ∆vap. H°