Chemical Bonds Chapter 22 Stability in Bonding Chapter

Chemical Bonds Chapter 22

Stability in Bonding Chapter 22, Section 1

Compounds & Chemical Formulas • Compounds: two or more elements that are chemically bonded together

Compounds & Chemical Formulas • Chemical formula: tells what elements a compound contains and the exact number of atoms of each element in a compound ▫ A combination of chemical symbols and subscripts (numbers written below the symbols) ▫ Each subscript indicates the exact number of atoms of the element it comes directly after ▫ If there are no subscripts, the compound contains only one atom of that element

Compounds & Chemical Formulas • Ex: H 2 O has 2 atoms of H and 1 atom of O • PRACTICE - How many atoms of each element are in the following compounds: 1. Si. O 2 2. C 12 H 22 O 11 3. Mg(OH)2 4. CH 3 COOH

Atomic Stability • All elements want to reach stability • An atom is considered chemically stable when its outermost energy level has the maximum number of electrons ▫ Ex: hydrogen and helium are stable with two electrons ▫ Ex: the outer energy levels of all other elements are stable when they contain eight electrons

Atomic Stability • Valence electrons: electrons in the outermost energy level of an atom • Groups 1 A, 2 A, 3 A, 4 A, 5 A, 6 A, 7 A, and 8 A indicate the number of valence electrons in the Group # before the “A” ▫ Ex: Elements in Group 3 A (Group 13) have 3 valence electrons • Recall that the Noble Gases all have full outer energy levels ▫ This indicates that elements in Group 8 A (Group 18) are stable

Electron Dot Structures of Some Group A Elements Group Period 1 2 3 4 1 A 2 A 3 A 4 A 5 A 6 A 7 A 8 A

Bonding • Atoms with partially stable outer energy levels can lose, gain, or share electrons to obtain stable outer energy levels • They do this by combining with other atoms that also have partially complete outer energy levels ▫ Ex: sodium, Na bonding with chlorine, Cl Na+ + Cl- Na. Cl

Octet Rule • Octet rule: when forming compounds, atoms want to fill their outer energy levels with a full set of valence electrons • An octet is a set of 8 • Atoms want to look like the Noble Gases

Ions • Ion: a charged particle because it has either more or fewer electrons than protons • Cation: positively charged ion ▫ Formed when an atom loses one or more valence electrons ▫ Typically formed by metals ▫ Ex: sodium ion Na+ • Anion: negatively charged ion ▫ Formed when an atom gains one or more valence electrons ▫ Typically formed by nonmetals ▫ Ex: chloride ion Cl-

Ionic Bonding Chapter 22, Section 2 – Part 1

Chemical Bonds & Ionic Compounds • Chemical bond: the force that holds atoms together in a compound • Compounds have different physical and chemical properties from those of the atoms that make up the compound • Ionic compound: a compound made of positively charged ions and negatively charged ions ▫ Ionic compounds are neutral because the total positive charge of the cations equals the total negative charge of the anions

Chemical Bonds & Ionic Compounds • Recall that all atoms want to have a full set of valence electrons like the Noble Gases • Ex: when sodium and chlorine react to form a compound, the sodium atom transfers its one valence electron to the chlorine atom that is lacking one valence electron to have a full set ▫ Sodium and chlorine can combine in a 1: 1 ratio so that both achieve stable octets!

Valence Electrons and Ions • Metals in Group 1 A tend to lose their 1 valence electron • Metals in Group 2 A tend to lose their 2 valence electrons • Metals and Metalloids in Group 3 A tend to lose their 3 valence electrons • Elements in Group 4 A do not tend to form ions • Nonmetals in Group 5 A tend to gain 3 valence electrons • Nonmetals in Group 6 A tend to gain 2 valence electrons • Nonmetals in Group 7 A tend to gain 1 valence electron

Modeling Ionic Compounds • Recall that a chemical formula shows the numbers of atoms of each element in a compound using subscripts • In order to determine how many atoms of each element are required to form a bond so that the compound is electrically neutral, we can use electron dot diagrams

• Ex: When calcium and chlorine combine ▫ For every 1 calcium atom, 2 chlorine atoms are required to “balance” the charges and make the compound neutral ▫ The chemical formula is thus: Ca. Cl 2

Modeling Ionic Compounds • Use electron dot diagrams to predict the formulas of the ionic compounds formed from the following elements: 1. Potassium and oxygen 2. Aluminum and bromine 3. Magnesium and nitrogen

Ionic Bonds • Ionic bond: the force of attraction between the opposite charges of ions in an ionic compound • Ionic bonds are formed when metals bond with nonmetals • Ionic bonds involve the TRANSFER of electrons to form compounds

Ionic Bonds • Ex: in the example before, calcium “loses” or “transfers” its 2 valence electrons to two chlorine atoms (1 to each) ▫ The calcium atom thus attains a charge of +2 ▫ Each chlorine atom thus attains a charge of -1 ▫ The result is a compound with a net charge of zero! • Because of the new charges attained by each atom, the ions stick together by electric forces ▫ Positives are attracted to negatives

Naming & Writing Formulas for Ionic Compounds

Cations • Lose electrons • Formed by metals • Groups 1 A, 2 A, and 3 A form charges equal to their group number

Group 1 A: Lose 1 electron to form 1+ ions H+ Li+ Na+ K+

Group 2 A: Lose 2 electrons to form 2+ ions Be 2+ Mg 2+ Ca 2+ Sr 2+ Ba 2+

Group 3 A: Lose 3 electrons to form 3+ ions B 3+ Al 3+ Ga 3+

Naming Cations • The names of Groups 1 A, 2 A, and 3 A are the same as the name of the metal, followed by the word ion or cation Ion Na+ Ca 2+ Al 3+ Name Sodium ion Calcium ion Aluminum ion

Anions • Gain electrons • Formed by nonmetals • Groups 5 A, 6 A, and 7 A form charges equal to: 18 – group number Ex: elements in Group 7 A 17 – 18 = -1 Elements in Group 7 A form anions with a -1 charge!

Group 14 (4 A): there aren’t any common ions for this group.

N 3 P 3 As 3 - Nitride Phosphide Arsenide Group 15 (5 A): Gain 3 electrons to form 3 - ions

O 2 - Oxide S 2 - Sulfide Se 2 - Selenide Group 16: Gain 2 electrons to form 2 - ions

F- Fluoride Br- Bromide Cl- Chloride I- Iodide Group 17: Gain 1 electron to form 1 ions

Group 18/ 8 A: Stable Noble gases - do not form ions!

Naming Anions • Anion names start with the root of the element name and end in –ide Ion N 3 P 3 O 2 - Name Nitride Phosphide Oxide

Practice Naming Ions: • Write the symbol for the ion formed, classify it as a cation or anion, and name the ion for each of the following: 1. Potassium K+ , cation, potassium ion 2. Iodine I-, anion, iodide 3. Sulfur S 2 -, anion, sulfide

Naming Ions formed by Transition Metals • Ions formed by transition metals follow a different rule: we will not focus on these yet…

Naming Ionic Compounds • Recall that an ionic compound is made up of a metal cation and a nonmetal anion • Always name the cation first, followed by the anion 1. Cation = name of the element Ex: Ca 2+ = calcium (ion) 2. Anion = root of element name + -ide Ex: Cl- = chloride So the name of the ionic compound formed is calcium chloride

Practice Naming Ionic Compounds 1. Ca. Br 2 ▫ Calcium bromide 2. The ionic compound formed by combining potassium and oxygen ▫ Potassium oxide 3. Be 3 N 2 ▫ Beryllium nitride

Writing Chemical Formulas for Ionic Compounds Steps to writing chemical formulas for ionic compounds: 1. Write the chemical symbols with charges for both ions ▫ Be sure to write the cation’s symbol first! 2. Add subscripts when necessary to balance the charges by using the CRISSCROSS method

Writing Chemical Formulas for Ionic Compounds Example: Aluminum sulfide 1. Write the formulas for the cation and anion, including charges ▫ Al 3+ S 2 - 2. Use subscripts to balance the charges using the CRISSCROSS method, if necessary ▫ ▫ Al 3+2 S 2 -3 Correct formula: Al 2 S 3

Practice Writing Chemical Formulas for Ionic Compounds 1. Sodium nitride ▫ Na 3 N 2. Magnesium oxide ▫ Mg. O 3. Gallium bromide ▫ Ga. Br 3

Transition Metals and Ionic Bonding

Binary Ionic Compounds • Binary compound: a compound made up of two elements ▫ Ex: Na. Cl, Mg. Br 2 • In order to write the correct chemical formula, you need to know the correct charges of each ion involved in the bond ▫ CRISSCROSS method

Oxidation Numbers • Recall that all elements in the same group have the same number of valence electrons (electrons in their outermost energy levels) • Thus, they must gain or lose the same number of electrons to achieve a full outer energy level in a bond • Oxidation number: the charge of an ion • For ionic compounds, the oxidation number is the same as the charge of the ion ▫ Ex: the sodium ion (Na+) has a charge of 1+ and an oxidation number of 1+

Transition Metals & Other Metals • Transition metals are the elements in the main body of the periodic table • These are known as Type II metals • These metals tend to form cations with different oxidation numbers • The charges of the ions formed are determined by the number of electrons lost • Ex: iron (Fe) can lose 2 electrons (Fe 2+) or 3 electrons (Fe 3+)

Transition Metals & Other Metals • Common metals (Type II metals) that have more than one oxidation number: Name Oxidation Number Copper (I) 1+ Copper (II) 2+ Iron (III) 3+ Chromium (II) 2+ Chromium (III) 3+ Lead (II) 2+ Lead (IV) 4+

Naming Compounds of Metals with Multiple Oxidation Numbers • Place a Roman Numeral in parentheses after the name of the element to indicate the charge (oxidation number) ▫ Ex: Fe 2+ is iron (II) • Write the chemical formula in the same way you do for other ionic compounds by crisscrossing the charges ▫ Ex: Fe 3+ bonded with O 2 Fe 2 O 3

Practice Naming! Name the ionic compounds formed when combining the following ions: 1. Cu 2+ and Br▫ Copper (II) bromide 2. Pb 4+ and S 2▫ Lead (IV) sulfide 3. Cr 3+ and N 3▫ Chromium (III) nitride

Practice Writing Chemical Formulas! Write the correct chemical formula for the ionic compound formed by combining the following ions: 1. Copper (I) and selenide ▫ Cu 2 Se 2. Iron (III) and oxide ▫ Fe 2 O 3 3. Lead (II) and phosphide ▫ Pb 3 P 2

Polyatomic Ions and Ionic Bonds

Polyatomic Ions • Not all ionic compounds are binary (include only two elements) • Some ionic compounds contain more than two elements because they include polyatomic ions • Polyatomic ion: a group of atoms that are chemically bonded together and have a positive or negative charge together ▫ Ex: C 2 H 3 O 2 - = acetate ▫ Ex: NH 4+ = ammonium • Look at your table of common polyatomic ions to help

Practice Identifying Polyatomic Ions Write the correct chemical formula and charge for each of the following polyatomic ions: 1. Hydroxide ▫ OH- 2. Phosphate ▫ PO 43 - 3. Nitrate ▫ NO 3 -

Naming Ionic Compounds with Polyatomic Ions 1. Write the name of the cation (positive ion) first 2. Then write the name of the anion (negative ion) • Ex: K 2 SO 4 ▫ Potassium sulfate

Writing Chemical Formulas for Ionic Compounds with Polyatomic Ions • Follow the rules for naming binary compounds by crisscrossing charges ▫ List cation first, followed by anion • Use parentheses around each polyatomic ion before adding the appropriate subscript • Ex: barium chlorate ▫ Ba(Cl. O 3)2

Practice Naming! Write the correct names for the following ionic compounds: 1. Sr(OH)2 ▫ Strontium hydroxide 2. Ca. SO 4 ▫ Calcium sulfate

Practice Writing Chemical Formulas! Write the correct chemical formulas for the following ionic compounds: 1. Beryllium nitrate ▫ Be(NO 3)2 2. Sodium carbonate ▫ Na 2 CO 3

Ionic Bonding Quiz Important Information Name Chemical Formula Type II Metals Chlorate Cl. O 3 - Titanium Chlorite Cl. O 2 - Nickel Cyanide CN- Lead Acetate CH 3 COO- Iron Phosphate PO 43 - Cobalt Sulfate SO 42 - Carbonate CO 3 2 - Chromium Gold

Molecular Compounds

Covalent vs. Ionic Bonds • Recall that an ionic bond involves the transfer of valence electrons between a metal and a nonmetal • In a covalent bond, valence electrons are shared between two nonmetals

Covalent Bonds • Atoms involved in a covalent bond still want to achieve a full outer energy level • By sharing valence electrons, all atoms involved can count the electrons toward achieving a complete octet (set of 8 valence electrons)

Molecules • Molecule: a neutral group of atoms joined together by covalent bonds • Molecular compound: a compound composed of molecules • Diatomic molecule: a molecule that contains two identical atoms ▫ Br 2, I 2, N 2, Cl 2, H 2, O 2, F 2

Single Covalent Bonds • Single covalent bond: two atoms held together by sharing one pair of (2) electrons • Ex: All halogens have seven valence electrons ▫ By sharing valence electrons, two of the same halogens end up with a complete octet

Single Covalent Bonds • Ex: Similar to the halogens, hydrogen exists as a diatomic molecule, sharing its 1 pair of electrons

Double Covalent Bonds • Double covalent bond: a bond that involves two shared pairs of electrons ▫ Ex: Carbon dioxide (CO 2)

Triple Covalent Bonds • Triple covalent bond: a bond formed by sharing three pairs of electrons ▫ Ex: Nitrogen (N 2)

Steps to Drawing Lewis Structures 1. Count the total number of valence electrons for the given molecule by adding each individual element’s valence electrons 2. Draw the “skeleton” of atoms in the molecule ▫ Hydrogens are always around a central atom 3. Join all the atoms using only single bonds

Steps to Drawing Lewis Structures 4. Add the electrons as dots to fulfill the octet rule around each element in the compound Remember that hydrogen requires only 2 electrons 5. If the total comes out to what you calculated in step 1, you’re done! If not, you may need to add double or triple bonds as necessary Remember that each atom still needs to satisfy the octet rule!

Exceptions to the Octet Rule • The octet rule cannot be satisfied in molecules whose total number of valence electrons is an odd number ▫ Ex: NO 2 contains 17 total valence electrons – each oxygen contributes 6 and the nitrogen contributes 5 • There also molecules in which an atom has less, or more, than a complete octet of valence electrons ▫ Ex: sulfur hexafluoride (SF 6)

Practice Drawing Lewis Structures 1. 2. 3. 4. SCl 2 CH 4 NH 3 H 2 O

Bond Polarity & Hydrates

Unequal Sharing • Electrons are NOT always shared equally between atoms in a covalent bond • The strength of attraction of each atom to its electrons is related to: 1. The size of the atom 2. The charge of the nucleus 3. The total number of electrons in the atom

Unequal Sharing • Electronegativity: the ability of an atom to attract electrons in a chemical bond ▫ High electronegativity means that it pulls the electrons toward it ▫ Electronegativity decreases as you move down a group, but increases as you move from left to right across a period ▫ Fluorine is the most electronegative atom on the periodic table

Bond Polarity • The unequal sharing of electrons can give one atom a slight negative charge (δ-) and another atom a slight positive charge (δ+) • The atom holding the electrons more closely is the more electronegative atom and has a slight negative charge (δ-)

Bond Polarity • Polar molecule: a molecule with a slightly positive end a slightly negative end, although the overall molecule is neutral ▫ Ex: the O-H bonds in a water molecule

Bond Polarity • Nonpolar molecule: a molecule in which electrons are shared equally ▫ Ex: diatomic molecules (identical atoms), or symmetric molecules

Difference 0 or less Non. Polar Difference greater than 0 but less than 2 Polar Difference greater than 2 Ionic Bond

Ionic vs. Covalent Bonds • Covalent bonds are strong, but the attractions between individual molecules are weak • Ionic bonds are strong because of the strong attraction between positive and negative ions ▫ Therefore, ionic compounds have high melting and boiling points compared to covalent compounds

Hydrates • Hydrate: a compound that has water chemically attached to its ions and written into its chemical formula • To write the chemical formula for a hydrate, write the formula for the compound, then place a dot followed by the number of water molecules ▫ Ex: calcium sulfate dihydrate is Ca. SO 4 • 2 H 2 O

Naming Molecular Compounds & Hydrates

Binary Molecular Compounds • Binary molecular compounds are made up of two nonmetals • Prefixes are used to describe the number of atoms of an element that are present in each molecular compound

Naming Molecular Compounds • To name a binary molecular compound: 1. Write the names of the elements in the order listed in the chemical formula 2. Use prefixes appropriately to indicate the number of each kind of atom Omit the prefix mono- for the first element in the compound 3. End the name of the second element with the suffix –ide

Practice Naming Molecular Compounds 1. CO ▫ Carbon monoxide 2. Cl 2 O 8 ▫ Dichlorine octoxide 3. PCl 3 ▫ Phosphorus trichloride

Writing Chemical Formulas for Binary Molecular Compounds • To write the correct chemical formula for a binary molecular compound: 1. Use the prefixes to tell you the subscript of each element in the formula 2. Write the correct symbols for the elements with the appropriate subscripts

Practice Writing Chemical Formulas for Molecular Compounds 1. Tetraphosphorous trisulfide ▫ P 4 S 3 2. Iodine heptafluoride ▫ IF 7

Naming Hydrates • Use the same prefixes to name hydrates • To name a hydrate: 1. Name the ionic compound as you normally would 2. Write the prefix to indicate the number of water molecules in the compound 3. Follow the prefix with the word hydrate

Writing Chemical Formulas for Hydrates • To write the formula for a hydrate: 1. Write the chemical formula for the ionic compound 2. Draw a dot 3. Write the number of waters that are in the compound

Practice Naming Hydrates Write the name of each of the following hydrates: 1. Co. Cl 2 • 6 H 2 O ▫ Cobalt (II) chloride hexahydrate 2. Cu. SO 4 • 5 H 2 O ▫ Copper (II) sulfate pentahydrate

Practice Writing Formulas for Hydrates Write the chemical formula for the following hydrates: 1. Iron (III) sulfate nonahydrate ▫ Fe 2(SO 4)3 • 9 H 2 O 2. Sodium carbonate decahydrate ▫ Na 2 CO 3 • 10 H 2 O