Types of Chemical Bonds

Types of Chemical Bonds • Covalent bond • Electrons shared • Both atoms have high affinity • Ionic bond • Atoms have very different affinities • Atom with greater affinity takes electron • Forms ions which are held together because opposite charges attract

Types of Chemical Bonds • Metallic bonds • Atoms have weak affinities & few valence electrons • Easily lost electrons are shared among many

Covalent Bonds • Atoms • Have similar electronegativities • Are only a few missing valence electrons • Are usually nonmetals • Share valence electrons

Covalent Bonds • Bonding pair • Electrons shared in the bond • Located between nuclei of atoms • Form area of negative charge, hold atoms together  electrostatic force

Diatomic Elements • Strong affinities, bond quickly • Exist in diatomic, pure form • Ex: H2, N2, O2, F2, Cl2, Br2, I2

Lewis Dot Structure • Shows bond between different atoms • - = 2 electrons • Nonbonding electrons shown as dots

Ionic Bonding • Large difference in electronegativities • Form ionic compounds • Cations and anions held by electrostatic attraction • No distinct molecules • Sum of charges is 0 • Formula unit: ratio of cations to anions  neutral • Forms crystal lattice (orderly arrangement of ions)

Metallic Bonding • Electron sea theory • Metals are a crystal lattice of metal cations in a “sea” of mobile electrons • Electrons shared among many • Delocalized Electrons • Mobile • Hold cations together

6C Properties of Compounds • Covalent Compounds • Form molecules • Held by weak forces, separated easily • Poor conductors, fairly low melting points • Wide array (range) of colors • EXCEPTION: Network covalent substances • Hardest known substances • Ex: diamonds, silicates, quartz • Hard, brittle crystals

Properties of Compounds • Ionic Compounds • Dense, brittle, hard solids because held by strong electrostatic forces • High melting points • Split or cleave along flat surfaces • Poor conductors because valence electrons are held tightly • Conduct electricity if soluble in water

Properties of Metals & Alloys • Due to delocalized electrons, metals are… • Conduct heat and electricity • Have luster (e- absorb & reemit light) • Malleable (e- act as glue for metal cations)

Orbitals & Valence Bond Theory • Orbitals: • Each holds 2 electrons • Areas where electrons may be found • In different sublevels (s,p,d,f) of energy levels • S – 1 orbital • P – 3 orbitals • D – 5 orbitals • F – 7 orbitals

Orbitals with unpaired electrons overlap  forms a bond • Called valence bond (a model that explains how covalent bonds form)

Sigma bonds (σ) • End to end overlap of spherical s orbitals • Region of high electron density forms on the “line” that connects the two nuclei • Strongest type of covalent bond • Why? Electrostatic forces only pull nuclei together • Ex: H2

Double bonds • End-to-end overlap (sigma bond) and side-to-side overlap (pi bond) • Pi bond: two lobes overlap • Bond is parallel to bond axis (above & below) • 2 regions but 1 orbital! • Double bond = 1 sigma & 1 pi bond

7B Molecular Geometry • VSEPR: valence shell electron pair repulsion theory • Focuses on location of highest electron densities around central atom • 1 area = Single bond, double bond, triple bond, or unbonded pair of electrons • Charge repulsion determines shape • Areas of negative charge repel to be as far apart as possible

# of regions & shapes • 4 regions • Tetrahedral, 109.5 degrees apart • Ex: CCl4 • Pyramidal also has 4 regions • Ex: NH3 • 3 regions • Trigonal planar • Y shaped on single plane • Ex: CH2O • 2 regions • Linear, 180 degrees apart • Ex: diatomic molecules, CO2

Other (4 Region) Shapes • Bent • Typically 109.5 degree angle but 104.5 degrees in water • Nonbonding electrons repel more than bonding pairs • Linear • Ex: HF

Orbital Hybridization • Valence bond theory fails to explain how C bonds (should form 2 bonds since 1s2s 2p) • Suggested explanation: • C forms 4 hybrid sp orbitals • Hybridization • Process by which new orbitals with equal energies are formed by combining orbitals with different energies

Why hybridize? • Forms larger lobes, which are able to overlap other orbitals more effectively & form stronger bonds (more stable)

Polar Covalent Bonds • Electrons unevenly shared • Creates partial charges • Affects chemical properties • Polarity shown by Lewis dot structure symbol (+ ) • Arrow points to more electroneg. atom • Remember: greater ΔEN = more polar bond 

Dipole Moment • Molecules A through D are shown below. Determine whether each one is polar or nonpolar. • A – Nonpolar • B – Polar • C – Polar • D – Nonpolar D A B C

If a molecule has polar bonds, will it be polar? • For asymetrical molecules: YES • One end is + and other is – • Shapes: bent & pyramidal (always polar) • For symetrical molecules: Maybe • Tetrahedral, trigonal planar, & linear shapes are… • Not polar if outermost atoms are the same • Polar if outermost atoms are different

Oxidation Numbers & Formulas • Oxidation Numbers (ON) • Represent number of electrons element must gain/lose to return to neutral • Negative ON = bonded atom gained e- • Positive ON = bonded atom lost e- • May refer to charge of ionic compound/ion

Oxidation Number Rules • #1: Free-element Rule • ON of pure elements (& covalently bonded elements) is zero • #2: Ion Rule • ON of monoatomic ion is equal to charge of ion

Oxidation Number Rules • #3: Zero-sum Rule • Sum of ON of all atoms in compound must add up to zero • Element with highest electronegativity often determines ON of other atoms

#4: Specific ON Rule • Alkali metals: +1 • Alkali earth metals: +2 • Hydrogen: usually +1 if bonded to nonmetal, -1 if bonded to metal • Forms metallic hydrides • Oxygen: -2 unless bonded to F (-1 in rare cases) • Halogens: -1 when bonded to metals • F always has -1 ON • *If one rule contradicts another, rule listed first should be followed

Finding Oxidation Numbers • Algebra can be used (because sum of neutral compound equals zero) • ON numbers used to write chemical formulas • For binary (2 element) ionic compounds, use criss-cross method to find chemical formula

List the elements (positive ON always first) Write the ON underneath Draw the criss-cross arrows Rewrite the formula showing subscripts Formulas should show simplest ratio Mg F +2 -1 Mg F +2 -1 MgF2 1:2 is simplest ratio Criss-Cross Method

Binary Covalent Compounds • Use Greek prefixes • Write least electronegative element first • Change ending of 2nd element to -ide • NOT acids • Mono only for 2nd element • Omit double vowels (mono-oxide becomes monoxide)

Binary Covalent Compounds • Application: CO2 P2S CO • Use Greek prefixes • Write least electronegative element first • Change ending of 2nd element to -ide • NOT acids • Mono only for 2nd element • Omit double vowels (mono-oxide becomes monoxide)

Binary Ionic Compounds • Greek prefixes not used • Positive ions keep original (parent) name • Negative ions have –ide ending • Name cation, then anion • If metal hydride (Metal + H), name metal 1st

Binary Ionic Compounds • Application: MgO Al2O3 MgH2 • Greek prefixes not used • Positive ions keep original (parent) name • Negative ions have –ide ending • Name cation, then anion • If metal hydride (Metal + H), name metal 1st

Ionic Compounds with Multiple Oxidation States • Use Stock System • Place Roman numeral after the metal to show its oxidation number • For Fe, Co, Cu, Mg, Pb, Sn • Ex: FeCl2 iron (II) oxide

Binary Acids • H as 1st element • Ex: hydrogen chloride HCl • When dissolved in water, add prefix “hydro-” & suffix “-ic acid” • Ex: HCl in water becomes hydrochloric acid

Chemical Equations • Show composition of substances (what & how much) Ca(HCO3)2 + Ca(OH)2  H2O + CaCO3 • Account for all atoms involved (law of mass conservation) • Balanced using coefficients (called balancing by introspection) • Is the formula balanced?

Balancing by Introspection • 1. Write the formula for reactants & products • 2. Check to see if equation is balanced • 3. Adjust coefficients to balance • 4. Make sure coefficients are in simplest whole-number ratio • Practice: Work through example problem (p. 200-201), then do SRQ 8C 2a,b,c • Take notes on “Special symbols & equations” section

eggs shoes Molar mass is the mass of 1 mole of in grams marbles atoms For any element atomic mass (amu) = molar mass (grams) 1 12C atom = 12.00 amu 1 mole 12C atoms = 6.022 x 1023 atoms = 12.00 g 1 mole 12C atoms = 12.00 g 12C 1 mole lithium atoms = 6.941 g of Li 3.2

1 mol K 6.022 x 1023 atoms K x x = 1 mol K 39.10 g K 6.022 x 10^23 6.022 x 10^23 6.022 x 10^23 amu amu amu How many atoms are in 0.551 g of potassium (K) ? 1 mol K = 39.10 g K 1 mol K = 6.022 x 1023 atoms K 0.551 g K 8.49 x 1021 atoms K 3.2

1S 32.07 amu 2O + 2 x 16.00 amu SO2 SO2 64.07 amu Molecular mass of a compound can be found by Adding together the molar masses of the elements (think of it as the total amu) *units: amu For any molecule molecular mass (amu) = molar mass (grams) 1 molecule SO2 = 64.07 amu molar mass = 1 mole SO2 = 64.07 g SO2  molecular mass 64.07 g SO2 3.3

Types of Formulas • Structural Formula • Shows the types of atoms • Exact composition • Arrangement of chemical bonds • Molecular Formula • Shows the types and numbers of atoms • More convenient • Does not show the shape, location or types of bonds

Types of Formulas • Empirical Formula • Tells what elements are present using the simplest whole number ratio • Formula units [ NaCl, CA(OH)2 ] are always empirical formulas • Example - Ethene • Molecular – C2H4 • Empirical – CH2

Percent composition • Shows what percentage of a compound’s total mass comes from each element • Based on the model, estimate what percent of the mass comes from H & what percent comes from O 3.5

Percent composition • Based on the formula, what percent of the mass comes from H and what percent comes from O? Total = O + H + H = 16.00 + 1.008 + 1.008 = 18.016 amu % Oxygen = 16.00 / 18.016 = 88.8% % Hydrogen = (2 x 1.008) / 18.016 = 11.2% 3.5

2 x (12.01 g) 6 x (1.008 g) 1 x (16.00 g) n x molar mass of element %C = %H = %O = x 100% = 52.14% x 100% = 13.13% x 100% = 34.73% x 100% 46.07 g 46.07 g 46.07 g molar mass of compound C2H6O Percent composition • Shows what percentage of a compound’s total mass comes from each element nis the number of moles of the element in 1 mole of the compound 52.14% + 13.13% + 34.73% = 100.0% 3.5

53.28 g 6.72 g %H = %O = x 100% = 11.2% x 100% = 88.80% 60.00 g 60.00 g Percent composition of an element in a compound A 60.0 g sample of water is decomposed into its elements of 53.28 g of O and 6.72 g of H. Find the percent composition. H2O 11.2 + 88.80% = 100.0% 3.5

128 g 91.96 g 64.14 g %O = %Na= %S = x 100% = 45.05% x 100% = 32.37% x 100% = 22.58% 284.1 g 284.1 g 284.1 g Percent composition of an element in a compound Find the percent composition of a substance with 91.96 g Na, 64.14 g S, and 128 g O. 32.37% + 22.58% + 45.05% = 100.0% 3.5

4 x 16.00 g 32.07 g 2 x 22.99 g %Na= %O = %S = x 100% = 22.58% x 100% = 45.05% x 100% = 32.37% 142.05 g 142.05 g 142.05 g Percent composition of an element in a compound Find the percent composition of Na2SO4. Na2SO4 32.37% + 22.58% + 45.05% = 100.0% 3.5

Percent Composition Elemental (sometimes called “laboratory”) Analysis • Listing the mass of each element present in 100 g of a compound • Can be used to find the empirical formula What is the chemical formula for a compound with The following elemental analysis? 49.5% C 5.2% H 28.8%N 16.5% O Steps: 1) convert % to g 2) convert to mol 3) divide by smallest number 4) round

C = H = N = O = 49.5% 5.2% 28.8% 16.5% Percent composition of an element in a compound In 100 g of a sample you would have 49.5 g of C In 100 g of a sample you would have 5.2 g of H In 100 g of a sample you would have 28.8 g of N In 100 g of a sample you would have 16.5 g of O

Types of Chemical Bonds