Chapter 2 Atoms Molecules and Ions Chemistry 1061

Chapter 2: Atoms, Molecules, and Ions Chemistry 1061: Principles of Chemistry I Andy Aspaas, Instructor

Atomic theory of matter • Dalton’s atomic theory – All matter composed of indivisible atoms • Atom: small particle, retains identity in reactions – Element: type of matter composed of only one kind of atom – Compound: type of matter composed of fixed proportion of 2 or more elements – Chemical reaction: rearrangement of atoms to give new chemical combinations

Deductions from Dalton’s Atomic Theory • Law of conservation of mass • Law of definite proportions – Example: 1. 00 g C reacted with oxygen – 2 compounds formed • 1. 3321 g O : 1. 00 g C • 2. 5542 g O : 1. 00 g C

Structure of the Atom • Dalton said atoms were indivisible, but experiments starting around 1900 showed that atoms themselves consist of particles – Nucleus: atom’s central core • Positively charged • Contains most of atom’s mass – Electron • Very light, negatively charged particles

Discovery of the electron • 1897, British physicist J. J. Thomson – Experiments showed atoms are not indivisible – Cathode ray tube – Calculated ratio of electron’s mass to its charge • 1909, American Robert Millikan – Charge on electron: 1. 602 x 10 -19 coulomb – Mass: 9. 109 x 10 -34 kg

Nuclear model of the atom • 1911, British physicist Ernest Rutherford – Gold foil experiment – Most alpha particles pass straight through gold foil – Most of atom = empty space – If a golf ball was the nucleus, the atom would be about 3 miles in diameter

Gold foil experiment

Nuclear isotopes • Nucleus comprised of 2 kinds or particles – Protons • Have positive charge equal in magnitude to electron’s negative charge • More than 1800 times more massive than electron – Neutrons • Mass nearly identical to proton, no charge

Atomic number and mass number • Atomic number (Z): total number of protons in the nucleus of an atom – New element definition: substance whose atoms all have the same atomic number • Mass number (A): total number of protons and neutrons in a nucleus • Nuclide: atom characterized by certain mass number and certain atomic number • Nuclide symbol: Z subscript, A superscript, element

Isotopes • Isotopes: atoms whose nuclei have the same atomic number but different mass numbers • Some elements have only one naturally occurring isotope (sodium-23) • Some have several naturally occurring isotopes: – Oxygen contains the following isotopes • 99. 769% oxygen-16 • 0. 037% oxygen-17 • 0. 204% oxygen-18

Atomic weights • Dalton: an atom of a certain element has a characteristic mass • But, naturally occurring elements may be a mixture of isotopes – Isotope percentages are generally constant • Dalton actually calculated average atomic masses – They were relative to the mass of hydrogen, the smallest element

Atomic mass units and atomic weight • Carbon-12 isotope chosen arbitrarily as standard – 12 atomic mass units (amu) • Atomic weight: average atomic mass for a naturally occuring element, expressed in amu – Appears on periodic table along with atomic number

Calculating atomic weight • Fractional abundance: fraction of a total number of atoms, which consists of a particular isotope • Isotopic mass is not exactly equal to mass number – Neon-20, mass = 19. 992 amu, abund = 0. 9051 – Neon-21, mass = 20. 994 amu, abund = 0. 0027 – Neon-22, mass = 21. 991 amu, abund = 0. 0922 • Multiply isotopic mass by fractional abundance, and sum total to get atomic weight • Gives a “weighted average”

Periodic table of the elements • 1869, Dimitri Mendeleev (and J. Lothar Meyer, independently) – Arranged elements in order of atomic number – Placed elements in horizontal rows so that vertical columns of elements with similar properties formed

Periods • Period: elements in any one horizontal row of the periodic table – First has only 2 – 2 nd and 3 rd have 8 – 4 th and 5 th have 18 – 6 th has 32 – 7 th is incomplete

Groups • Group: elements in any one vertical column of the periodic table – Are numbered with Roman numeral and letter – (or just numbered with a number) • Main-group elements (labeled with A) • Transition elements (labeled with B) • Inner transition elements (shown below rest of table)

Groups • Elements in a single group have similar properties – For example: • IA: alkali metals - all react vigorously with water (except hydrogen) • VIIA: halogens - all react vigorously with sodium

Metals, Nonmetals, Metalloids • Metal: substances or mixtures that… – Have a characteristic shine – Generally are good conductors of heat and electricity – Are usually malleable (can be hammered into sheets) – Are usually ductile (can be drawn into wire) – Metallic elements are solid at room temp. , except for mercury

Metals, nonmetals, metalloids • Nonmetal: element that does not exhibit characteristics of a metal – Most are gases – Solid nonmetals are usually brittle and hard – Bromine is only liquid nonmetal • Metalloid (or semimetal) - both metallic and nonmetallic properties – Many are semiconductors (silicon, germanium) – Don’t conduct when pure, do when doped or at very high temperature

Chemical formulas • Chemical formula: atomic symbols with numeric subscripts that show relative proportions of atoms of different elements in a substance • Al 2 O 3: ratio of Al atoms to O atoms is 2 : 3 • No subscript: ratio of 1 (Na. Cl, 1: 1)

Molecular substances • Molecule: group of atoms that are chemically bonded together - tightly connected by attractive forces • Molecular substance: composed of all the same molecules • Molecular formula: shows number of different atoms of an element in a molecule • Structural formula: shows how atoms are bonded together, a line indicates a chemical bond

Molecular substances • Some elements are molecular substances – Cl 2, O 2, N 2, S 8 • Some exist as individual atoms – He, Ne • Some exist as a very large but indefinite number of atoms bonded together –C • Polymer: extremely large molecules made of small molecules repeatedly bonded together (monomers) – Natural (wool, cotton), synthetic (plastics, Nylon, polyester, Kevlar)

Ionic substances • Unlike molecular substances, ionic substances are formed from charged atoms or groups of atoms called ions – Metal atoms tend to lose electrons: result is a positive charge (cation) – Nonmetals tend to gain electrons: result is a negative charge (anion) • Sodium loses one electron to become Na+ • Calcium loses two electrons to become Ca 2+ • May consist of a group of bonded atoms that have a deficiency or surplus of electrons (SO 42+)

Ionic compounds • Ionic compounds are composed of cations and anions • Ionic formula: given by smallest possible integer number of different ions in substance – Charges are omitted – Ionic compounds are uncharged as a whole – Ex. Na+ and Cl- = Na. Cl – Fe 3+ and SO 42 - = Fe 2(SO 4)3

Naming simple compounds • Chemical nomenclature: systematic naming of chemical compounds • Organic compounds: carbon-containing molecular substances • Inorganic compounds: composed of elements other than carbon (except CO, CO 2, cyanides, and carbonates)

Ionic compounds • Named given by name of cation followed by name of anion – Potassium sulfate • Monatomic ion: formed from a single atom

Predicting a monatomic ion’s charge • Most main-group metallic ions have one monatomic cation with charge equal to group number – Aluminum, Group IIIA, Al 3+ • Many high-atomic-number metals have a cation equal to group number AND one equal to the group number minus 2 – Thallium, Group IIIA, Tl 3+, Tl+ • Most transition elements can have several monatomic cations. +2 is very common • A nonmetal main-group ion forms an anion – Charge = 8 - group # – Oxygen, Group VIA, O 2 -

Naming monatomic ions • Monatomic cations are named after element if there’s only one possible cation for that element • If more than one possible cation, use charge as roman numeral after name – Fe 2+ = iron(II), Fe 3+ = iron(III) – Or use suffixes with latin name (-ous for lower charge, -ic for higher) • Cu 2+ = cuprous, Cu 3+ = cupric • Monatomic anions use -ide suffix (bromide, sulfide)

Polyatomic ions • Polyatomic ion: ion consisting of 2 or more atoms chemically bonded together, which carry a net electric charge • Oxoanions: central element + some oxygens – Suffixes indicate relative amount of oxygens – Fewer oxygens: -ite – More oxygens: -ate – Ex. SO 32 - is sulfite, SO 42 - is sulfate

Polyatomic anions • If more than 2 oxoanions, use prefixes – hypo- and per • Cl. O- hypochlorite ion • Cl. O 2 - chlorite ion • Cl. O 3 - chlorate ion • Cl. O 4 - perchlorate ion

Binary molecular compounds • Name first element with its exact name • Suffix second element with -ide • Use Greek prefixes to indicate number • Don’t use mono unless there are more than one possible amount – Ex. CO vs CO 2

Acids and corresponding anions • Acids have H+ as the cation • Oxoacids contain an oxoanion – -ate becomes -ic – -ite becomes -ous – End name with “acid” • Binary compounds of hydrogen and nonmetal – Prefix hydro– Suffix -ic acid

Hydrates • Simply indicate how many water molecules are present by using the Greek numerical prefix and hydrate at the end • Mg. SO 4 • 7 H 2 O = magnesium sulfate heptahydrate

Writing chemical reactions • Reactant ----> Product • Arrow means “reacts to form” or “yields” • Useful to indicate state or phase – (g) = gas, (l) = liquid, (s) = solid, – (aq) = aqueous solution • Use a coefficient to indicate relative number of particles involved • ∆ over arrow means heat is applied • A compound written over the arrow is usually a catalyst

Balancing chemical equations • Mass must be conserved so use coefficients to make sure the same number of each atom occurs on each side of the equation • Start by balancing atoms for elements that occur in only one substance on each side – Ex: H 3 PO 3 ---> H 3 PO 4 + PH 3 – Start by balancing oxygen – Practice!