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8. Bonding: general concepts

Photoelectron spectroscopy

a machine tells you the electron configuration of an atom

relative heights show how many electrons

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Types of bonds

ionic bonds

take electrons from each other to make ions

ions attract

metal and nonmetal

Coulomb’s law:

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E: energy (J)

k: constant

Q: charge of ions

r: length of bond

when a bond is made, energy is released (exothermic)

the higher the charge, the stronger the bond

the shorter the bond, the stronger the bond

find with atomic radius

covalent bonds

share electrons

nonpolar: share equally

polar: share unequally

nonmetal and nonmetal or semimetal

the shorter the bond, the stronger the bond

Electronegativity

electronegativity: how much an atom pulls electrons towards its nucleus

across: increases

down: decreases

fluorine is strongest (4.0)

many protons

small atomic radius

shielding: electrons block positive charge

the more energy levels you have, the stronger the shielding

example: H—H, O—H, Cl—H, S—H, F—H

put in order of increasing polarity

H: 2.1

S: 2.5

Cl: 3.0

O: 3.5

F: 4.0

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H—H: |2.1 - 2.1| = 0

O—H: |3.5 - 2.1| = 1.4

Cl—H: |3.0 - 2.1| = 0.9

S—H: |2.5 - 2.1| = 0.4

F—H: |4.0 - 2.1| =1.9

H—H < S—H < Cl—H < O—H < F—H

≤ 0.3

nonpolar

0.3 - 1.7

polar

≥ 1.7

ionic

There are no rows in this table

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Ionic compounds

valence electrons

1s²2s²2p⁶3s¹

1 valence electron

1s²2s²2p⁶3s²3p⁵

7 valence electrons

1s²2s²2p⁶3s²3p⁶4s²3d¹⁰4p⁵

7 valence electrons

octet rule: atoms want to have 8 valence electrons to be stable (s + p)

noble gases naturally have 8 valence electrons so they are not very reactive

hydrogen and helium: duet rule (two electrons)

form ions by taking/giving electrons

e.g. Na

lose one valence electron for octet rule

become Na⁺

1s²2s²2p⁶

e.g. Cl

add one valence electron for octet rule

become Cl⁻

1s²2s²2p⁶3s²3p⁶

combine ions

Na⁺ + Cl⁻ = NaCl

Fe³⁺ + S²⁻ = Fe₂S₃

Covalent compounds and dot diagrams

Gilbert Lewis: dot structures

1 dot = 1 valence electron

the one with more lone pairs goes in the middle

add atoms to fulfill octet/duet rule

types of bonds

single bond: two dots or one line

double bond: four dots or two lines

triple bond: six dots or three lines

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polyatomic dot diagrams

charges

positive: fewer electrons

negative: more electrons

put brackets around the diagram and add charge

exceptions to octet rule

H, He: 2 electrons

B: 6 electrons

everything from Period 3 downwards can have more than 8 electrons

at energy level 3, electrons start to use the d orbitals

isomer: compounds with same formula but different structures

resonance: dot diagrams with different bonding

e.g. NO₃⁻

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the bond for NO₃⁻ is an average of the different bonds (4/3 bond)

resonance: all the bonds are the same length

Covalent bond energies

the shorter the bond, the more energy it takes to break it

breaking bonds: endothermic

making bonds: exothermic

ΔH = energy to break bonds - energy to form bonds

reactants - products (opposite of usual!)

example:

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bond energies

C—H: 314 kJ/mol

Cl—Cl: 239 kJ/mol

F—F: 154 kJ/mol

C—F: 485 kJ/mol

C—Cl: 339 kJ/mol

H—F: 565 kJ/mol

H—Cl: 427 kJ/mol

draw Lewis dot structures for equation

multiply bond energies by number of bonds

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ΔH = -1194 kJ

Formal charge

formal charge: using math to prove dot diagrams correct

number of electrons assigned - number of electrons used in the compound

goal: as many atoms to have a formal charge as close to 0 as possible

use to check whether the Lewis structure works (type of bond, etc.)

example: carbon in CF₄

dots assigned: 4

dots used: 4

formal charge: 0

example: nitrogen in N₂

dots assigned: 5

dots used: 5

formal charge: 0

example: SO₄²⁻

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S: 6 - 6 = 0

O (single): 6 - 6 = 0

O (double): 6 - 7 = -1

two 1- charges → 2- charge overall

example: XeO₃

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Shapes of compounds

ionic: crystalline

e.g. NaCl

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covalent: VSEPR theory (valence shell electron pair repulsion)

electron pairs want to be as far apart from each other as possible

shape

number of atoms bonded to central atom

number of unshared pairs on central atom

name

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linear (180°)

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trigonal planar (120°)

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tetrahedral (109.5°)

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trigonal bipyramidal

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octahedral

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bent

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trigonal pyramidal

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T-shaped

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seesaw

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square planar

There are no rows in this table

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Photoelectron spectroscopy

Types of bonds

Electronegativity

Ionic compounds

Covalent compounds and dot diagrams

Covalent bond energies

Formal charge

Shapes of compounds

Gallery

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