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ME Flashcards

1
Q

Mixtures

A

contain two or more elements or compounds. The composition of a
mixture is variable and the chemical properties of the mixture are those of
the components.

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2
Q

Any sample of material will be

A
  1. A mixture

2. A pure substance: A compound or an element

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3
Q

Elements

A

substances that cannot be broken down into simpler

materials by chemical means

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4
Q

Compounds

A

chemical combinations of two or more elements

in definite ratio by mass

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5
Q

limiting reagent

A

reagent present in the lowest
(molar) amount relative to the products of the reaction and
which therefore determines (i.e. limits) the yield of product.

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6
Q

Law of Constant Composition

A

A given pure compound always contains the

same elements in the same proportions of mass

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7
Q

atom

A

smallest particle of an element
that retains the characteristic chemical
properties of that element

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8
Q

1 mole (avogadros number)

A

6.023x10^23

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9
Q

molar volume stp and rtp

A

stp: 22.4L
rtp: 24L

Revise what standard T, standard P, and reg T and P are

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10
Q

molecule

A

smallest independent
neutral particle of an element or
compound capable of independent
existence in the gas phase.

eg. HCl

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11
Q

diff between compounds and molecules

A

Compounds are molecules, they are made up of diff elements chemically combined. Molecules like H2 however are not compounds as they are not diff elements combined

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12
Q

specifying the entity

A

instead of “1 mole of hydrogen”, specify:

  • 1 mole of H atoms
    or
  • 1 mole of H2 molecules
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13
Q

Molar Mass

A

mass per mole of its

entities (atoms, molecules) and has units of g/mol.

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14
Q

empirical formula

A

simplest
whole number ratio of the atoms combining to make up the
pure substance

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15
Q

molecular formula

A

gives the actual numbers of each
kind of atom present in a single molecule of any molecular
substance

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16
Q

structural formula

A

molecular formula rewritten to
give some further information about the structure of the
molecule

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17
Q

reagents and products equation

A

Reagents -> Products

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18
Q

balanced chemical equation

A

tells us the relative molar ratio of reagents consumed and products produced during a chemical
reaction.

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19
Q

Chemical equation gives no info on

A
  • way reaction takes place at molecular/atomic level (eg. does not imply that 2 molecules of x reacts with 1 molecule of y)
  • rate of reaction
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20
Q

number of moles formulas

A

n = m/mr

n = vm/1000
v in cm^3

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21
Q

combustion of hydrocarbons

A

involves burning the
material in oxygen resulting in the formation of H2O and CO2 and
the release of heat

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22
Q

standard temperature

A

0°C / 273.15K

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23
Q

Standard pressure

A

1 atm / 760 Torr / 101.3kPa

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24
Q

Molarity

A

moles per litre

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25
Q

Dalton’s atomic theory

A
  • all matter consists of atoms
  • atom definition
  • atom of one element cant be converted to another elem
  • atoms of same element have same mass + chem properties
  • during chem reactions, atoms conserved
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26
Q

Discovery of electron - cathode ray tube observations + conclusions

A

1- ray bends in magnetic field: consists of charged particles

2- ray bends towards pos plate: consists negative particles

3- ray identical or any cathode: particles found in all matter

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27
Q

who named the electron

A

George Johnstone Stoney

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28
Q

J.J Thomson

A

measured ratio of mass of cathode ray particle to its charge

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29
Q

Robert Milikan

A

-measured charge of electron in 1909

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30
Q

How milikan measured charge of electron

A
  • X-rays used to knock electrons from gas molecules
  • Electrons stuck to tiny oil droplets
  • Adjusting applied electric field, drop could be slowed + suspended
  • allowed total charge to be measured
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31
Q

Mass of electron formula

A

Mass of electron = (mass/charge ratio) x charge

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32
Q

Thomson’s Plum Pudding Model

A

spherical atom of diffuse positively charged matter with electrons
embedded in it

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33
Q

Rutherford

A

proposed that positive particles, called protons, were in the nucleus

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34
Q

Rutherfords observations

A

•atom is space mostly occupied by electrons
• In centre is a tiny region named the nucleus, contains all of + charge and majority of mass of
atom.

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35
Q

Chadwick

A

discovered neutron, uncharged particle that also lies in nucleus

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36
Q

atomic number (Z)

A

of an element equals the number of protons in
the nucleus.

neutral atom- also no. of electrons

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37
Q

Mass Number (A)

A

of an element equals the total number of

protons and neutrons in the nucleus.

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38
Q

Isotopes

A

of an element are atoms that have different numbers of

neutrons and therefore different mass numbers

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39
Q

Numbers on each element n periodic table

A

In general:

Top: atomic number
Bottom: Atomic Mass

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40
Q

How masses of atoms are measured + unit

A

-measured relative to the mass of an atomic
standard
-standard is the ¹²C atom (mass is defined as exactly 12
atomic mass units)

unit: atomic mass unit (amu) or Dalton

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41
Q

the Dalton/amu

A

1/12 mass of a ¹²C atom

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42
Q

1 amu =

A

Has an absoute mass of 1.66054x10⁻²⁴ g

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43
Q

atomic mass and molar mass

A

for each element will have the same
numerical value but different units

Atomic mass: amu/Dalton
Molar mass: g/mol

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44
Q

Mass spectrometry

A

ul technique for measuring the mass and abundance of charged

particles from their mass/charge ratio (m/e)

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45
Q

How mass spectrometer works

A

•High energy electrons bombard sample, creating positively charged
particles
• Attracted towards series of negatively charged plates
• Particles paths bent by magnetic field, separates them by
m/e
• @ End of magnetic region, particles strike detector which
counts their relative positions and abundance

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46
Q

atomic mass

A

the average of its naturally occurring

isotopes weighted according to their abundance.

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47
Q

Calculate the atomic mass of C given that 12C and 13C have masses of 12
amu and 13.0034 amu.
13C constitutes 1.1% and the rest is 12C

A

Atomic mass = (12 x 0.989)+(13.0034 x 0.011) = 12.011 amu

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48
Q

electronic structure

A

describes the arrangement of electrons around

the nucleus

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49
Q

frequency

A

the number of complete
waves that pass a point per second (Units: 1/s
or Hz)

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50
Q

wavelength

A

distance between any
point on a wave and the corresponding point
on the next crest (or trough). The distance a
wave travels during one cycle (Units: m)

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51
Q

speed

A

distance a wave moves per unit time

Units: m/s

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52
Q

speed of light

A

3.00 x 10^8 ms-1

given in exam

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53
Q

speed of electromagnetic radiation formula

A

c = vλ

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54
Q

blackbody radiation

A

As a solid object is heated it gives off light. As
the temperature is changed the intensity and wavelength of the
emitted light changes –

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55
Q

Planck’s equation - energy of radiation

A

E = nhv

E = energy of radiation
v = frequency
h = Planck's constant
n = a positive integer (A quantum number)
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56
Q

smallest possible energy change

A

occurs when n = 1

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57
Q

change in energy formula

A

ΔE = hν = (h c)/ λ

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58
Q

photoelectric effect

A

When monochromatic light
of a specific frequency shines on a metal plate a
current flows.

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59
Q

threshold frequency

A

below this freq no current flows, diff for diff metals

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60
Q

Einsten’s photon equation

A

Eₚₕₒₜₒₙ= hν = ΔEₐₜₒₘ

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61
Q

gas passes thru slit and reflected by prism

A
  • does not create continuous spectrum

- line spectrum seen

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62
Q

line spectrum

A

consists of a series of lines at specific frequencies

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63
Q

spacing between lines and wavelength and velocity

A
  • Spacing decreases when λ decreases or v/E increases

- Spacing increases when λ increases or v/E decreases

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64
Q

Lyman series

A
  • shortest λ

- in UV

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65
Q

Balmer series

A

-visible

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66
Q

Ryberg Equation

A

𝝀 = 𝑹(𝟏/𝒏₁² − 𝟏/𝒏₂²)

n₁ = original energy level of electron
n₂ = new energy level
R = ryberg constant
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67
Q

Ryberg Equation for electrons that are removed

A

n₂ equals infinity when electrons are removed,

1/infinity = 0

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68
Q

n and series

A

n₁ = 1; UV series (Lyman)

n₁ = 2; visible series (Balmer),

n₁ = 3; IR series

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69
Q

Bohr and Rutherford’smodel: Energy of a photon

A

Eₚₕₒₜₒₙ = ΔEₐₜₒₘ = E𝒻ᵢₙₐₗ– Eₛₜₐᵣₜᵢₙ𝓰 = hν

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70
Q

Quantum number (n)

A
  • defines energy of electron
  • lower n implies its closer to nucleus - more energy to remove
  • for n=1, electron is in orbit closest to nucleus, for H atom: ground state
  • for n>1: excited state
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71
Q

Equation for calculating energy levels of an atom

A

𝑬 = −𝟐.𝟏𝟖×𝟏𝟎⁻¹⁸𝑱(𝒁²/n²)

where z is charge on nucleus (1 for H)

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72
Q

Change in energy when electron moves between two levels fequation

A

∆𝑬 = 𝑬𝒻ᵢₙₐₗ– Eₛₜₐᵣₜᵢₙ𝓰 = −𝟐.𝟏𝟖×𝟏𝟎⁻¹⁸𝑱(1/n𝒻ᵢₙₐₗ² - 1/ₛₜₐᵣₜᵢₙ𝓰²) = −𝟐.𝟏𝟖×𝟏𝟎⁻¹⁸𝑱(𝟏/𝒏₁² − 𝟏/𝒏₂²)

once we know the E diff between levels, we can use this to find the wavelength for the associated line

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73
Q

−𝟐.𝟏𝟖×𝟏𝟎⁻¹⁸𝑱

A

a variation of Ryberg constant

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74
Q

Equation for wavelength of any particle mass, m moving at speed, u
(de Broglie)

A

𝝀 = 𝒉/𝒎u

m = mass
u = speed
h = planck's constant
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75
Q

de Broglie’s equation, but with momentum

A

𝝀 = 𝒉/𝒎u = h/p

where p = momentum as momentum is a product of mass and speed

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76
Q

photons and momentum

A

-higher energy photons, with a shorter 𝝀, have greater momentum

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77
Q

Heisenberg’s uncertainty principle

A

it is impossible to know the position

and momentum of a particle simultaneously

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78
Q

Heisenberg’s Uncertainty Principle Equation

A

∆𝒙.∆𝒑 ≥𝒉/𝟒π

∆𝒙 = uncertainty in position
∆𝒑 = uncertainty in momentum
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79
Q

schrodinger’s equation

A

Check slides 5 lol

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80
Q

simplest form of schrodinger’s equation

A

ĤΨ = EΨ

E = energy of atom
Ψ = wave function/atomic orbital
Ĥ = the Hamiltonian,
Ψ = function that describes all accessible info about an electron in corresponding orbital
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81
Q

orbital - schrodinger

A

a mathematical function that describes an electron’s matter wave

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82
Q

Ψ²

A

the probability density

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83
Q

what do we know from schrodinger’s equation

A

we can know where an electron probably is

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84
Q

probability contour

A

encloses a volume in which we are likely to find an

electron within a certain probability

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85
Q

Three quantum numbers

A
  • Principal quantum number (n)
  • Angular momentum quantum number (l)
  • Magnetic quantum number (mₗ)
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86
Q

principal quantum number

A

n = 1, 2, 3, etc

-orbital size and energy level

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87
Q

angular momentum quantum number (l)

A

l = 0, 1, … n-1

  • orbital shape
  • stands for s, p, d, f etc
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88
Q

magnetic quantum number

A

mₗ) = -l, -l+1, … 0 … l-1, l

-orbital orientation
l=0 (s) 2e- in one orbital

l=1 (p) 2e- in each of three sub orbitals

l=2 (d) 2e- in each of 5 sub orbitals

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89
Q

spin (mₛ)

A

mₛ = -1/2 or +1/2

-spins in any single sub-orbital must be paired

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90
Q

outermost occupied orbital

A

said to be in the valence shell

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91
Q

Pauli exclusion principle

A

No two electrons in the same atom can have the

same set of four quantum numbers

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92
Q

orbitals and electrons

A

0 electrons : empty (or unoccupied orbital)

1 electron : occupied

2 electrons : occupied - Filled orbital

All orbitals have Ψ = 0 at r = ∞

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93
Q

angular lodes

A

-Planar nodes which pass through the nucleus
- Each orbital will have l angular nodes
- When l = 0 (s orbitals), no nodes at nucleus
- s orbitals: only ones that have significant electron density at
nucleus

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94
Q

radial nodes

A
  • spherical surfaces

- each orbital will have (m-1-l) radial nodes

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95
Q

s orbital

A
  • spherical shape with nucleus at centre
  • l = 0, so ml = 0
  • only one orientation
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96
Q

p orbital

A
  • dumbbell shape - two lobes of high electron probability with node at
    nucleus
  • l = 1, so ml = -1, 0, 1
  • 3 orientations
  • mutually perpendicular along axes - pₓ, pᵧ, p𝓏
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97
Q

d orbital

A
  • l = 2, so ml = -2, -1, 0, 1, 2
  • 5 orientations
  • Four of these have four lobes – like cloverleaf (dxy, dxz, dyz, dx²-y²)
  • Fifth (dz²)has 2 lobes along z axis and a donut shaped region along the x-y
    plane
98
Q

As Z increases

A
  • lowers the sublevel energy (to a more negative energy)
  • harder to remove an electron – stabilises atom
  • orbitals get smaller
99
Q

For ions with only one electron

A
  • results same for h but larger Z

- multiply energy for H by Z²

100
Q

shielding

A

-electrons subject to electron-electron
repulsion as well as attraction to the nucleus

-repulsions will counteract the nuclear attraction

  • shielding:
  • pushes them further from nucleus
  • makes them easier to remove
101
Q

shielding and nuclear charge

A

shielding means that the full nuclear charge Z is reduced to an
effective nuclear charge Zₑ𝒻𝒻

-Electrons at a higher energy level experience a much lower Zₑ𝒻𝒻

102
Q

Why does 2s fill before 2p?

A

From radial distribution functions:

  • 2p on average slightly closer to nucleus
  • small part of 2s peaks (or penetrates) very close to nucleus

This has two main effects:

  • Dec shielding of 2s electrons by 1s electrons
  • Inc nuclear attraction for a 2s electron over 2p
103
Q

pauli principle

A

no more than two electrons per orbital

104
Q

aufbau principle

A

in the ground state of an atom or ion, electrons fill atomic orbitals of the lowest available energy levels before occupying higher levels

105
Q

hund’s rule

A

When orbitals of equal energy are available the electron
configuration of lowest energy has the maximum number of unpaired
electrons with parallel spins

106
Q

condensed electron configs

A

electron configurations can be written with the element
of the previous noble gas in brackets

eg. Si (Z = 14) = [Ne]3s23p2

107
Q

writing electron configs

A

remember aufbau principle, eg. 4s filled before 3d

108
Q

Periodic table groups and blocks

A

groups 1 & 2: s block

transition (valence) elements: d block

groups 3-8: p block

bottom 2 periods: f block

109
Q

cation

A

+

loss of electrons

110
Q

anion

A

gain of electrons

111
Q

Transition metals and electronic configurations

A

-in a cation, 3d is lower in energy than 4s, so if removing electrons, remove 4s electrons before 3d

and similarly for 5s/4d and 6s/5d

112
Q

spin quantum numbers

A

first electron to fill a particular orbital given spin quantum number mₛ = +1/2

second: mₛ = -1/2

113
Q

periodic table periods

A

the period number is the n value of the highest energy level

114
Q

alkali metals

A

GROUP 1A

  • outer electron config ns¹
  • highly reactive
  • lose outer reactions when form ionic comps
  • soft metals w/ low melt points
115
Q

Alkaline Earth Metals

A

GROUP 2A

  • Outer electron config ns²
  • greater Zₑ𝒻𝒻 than 1A
  • lose outer electrons to form cations when form ionic comps
  • higher melt points than 1A, harder + denser
116
Q

Halogens

A

GROUP 7A

  • highly reactive non metals
  • occur as diatomic molecules
  • form ionic comps w/ metals
117
Q

Noble Gases

A

8A

  • filled energy level
  • unreactive monatomic gases
118
Q

Group 3A-5A

A
  • wide variety of phys + chem properties

- phys properties depend on type of bonding in element

119
Q

The Oxygen Family

A

GROUP 6A

  • outer electron config ns²np⁴
  • phys properties differ greatly moving down
120
Q

The elements of the oxygen family

A

O: diatomic gas

S: polyatomic nonmetal

Se: nonmetal

Te: semi metal

121
Q

Periodic table - moving down group

A

n dominates as there is one more level of inner
electrons screening Z
→ atomic size generally increases

122
Q

Periodic table - moving across period (left to right)

A

Zₑ𝒻𝒻 dominates as the changes in
shielding are less significant (electrons added to same level) and Zₑ𝒻𝒻 gets
larger
→ atomic size generally decreases

123
Q

Anion effect on size of atom

A

adding an electron increases repulsions between all electrons and
so increases size – anion larger than corresponding atom

124
Q

cation effect on size of atom

A

cation smaller than corresponding atom

125
Q

ionisation energy

A

energy required to completely remove an electron
from an atom or ion

typical unit: kJ/mol

126
Q

ionisation enegy notes

A

-always positive

127
Q

first ionisation energy

A

(IEₗ) given by:

Atom -> ion⁺ + e⁻

ΔE = IEₗ > 0

128
Q

trends in periodic table

A

revise

129
Q

Trends in periodic table: moving down group

A 
• n gets bigger, atom gets bigger
• Distance from outer electron to
nucleus increases
• Attraction between nucleus +
outer electron lessens
• Electron is easier to remove (IE
decreases)
130
Q

Trends in periodic table: moving across period

A 
• Zₑ𝒻𝒻 increases + atomic size
decreases
• Attraction between nucleus +
outer electron increases
• Electron is harder to remove (IE
increases)
131
Q

electron affinity

A

energy released when an atom or ion gains an
electron

typically kJ/mol

132
Q

electron affinity notes

A
  • in most cases it is negative, as first electron attracted to nucleus
  • second will always be positive as energy must be absorbed to overcome electrostatic repulsion
133
Q

first electron affinity

A

EA₁ given by:

Atom + e⁻ -> ion⁻

ΔE = EA₁

134
Q

Reactive metals, groups 1A and 2A (e.g. Na)

A
  • Low IE and slightly negative EA
  • Lose electrons easily – attract them weakly
  • Likely to form positive ions
135
Q

Reactive non metals, group 7A (e.g. Cl)

A
  • Have high IE and highly negative EA
  • Do not lose electrons easily – attract them strongly
  • Likely to form negative ions
136
Q

Noble gases

A
  • Very high IE and slightly positive EA

* Do not tend to lose or gain electrons

137
Q

Electronegativity (χ)

A

measure of the electron-attracting power

of its atoms when they are part of a compound

138
Q

what electroneg means

A

• high electronegativity means a strong electron attracting power
• low electronegativity means at atom is more likely to lose electrons when forming a
compound

139
Q

Comparing χ and EA

A

χ: bonded atom attracting electrons

EA: Atom gaining electron to form an ion

140
Q

Electroneg trends - moving down group

A
  • n gets bigger, so atom gets bigger
  • The larger the atom the further it is from shared electrons
  • Electronegativity gets smaller
141
Q

Electroneg trends - moving across period

A
  • Zₑ𝒻𝒻 increases and atomic size decreases
  • The smaller the atom, the closer it is to shared electrons
  • Electronegativity gets larger
142
Q

Pseudo Noble gas configuration

A

(n – 1)d¹⁰ configurations

  • Metal empties its highest energy level (ns and np)
  • Attains stable configuration with empty ns and np and filled (n - 1)d
143
Q

Inert Pair Configuration

A
  • Metal loses just its np electrons
  • Attains stable configuration with empty np and filled ns and (n – 1)d
  • Retained ns – inert pair
144
Q

compound formation

A

reflects the tendency of atoms to achieve a noble
gas configuration – or a configuration with two electrons in the outermost
occupied s orbital and six electrons in the outermost occupied p sublevel

145
Q

octet rule

A

each atom in a compound will tend to have a total of eight

electrons in the valence s and p orbitals

146
Q

ionic bonding

A

When one or more electrons transfer from one atom to
another resulting in the formation of positively charged ions (cations) and
negatively charged ions (anions)

147
Q

covalent bonding

A

Sharing of pairs of electrons between atoms

148
Q

binary ionic compound

A
  • simplest type of ionic comp
  • typically forms when a metal reacts with non metal
  • metal loses one or more electron → cation
  • non metal gains one or more electrons → anion
149
Q

binary ionic compounds

A

-esulting compound is a crystalline solid, held together by the
electrostatic forces between the oppositely charged ions
-lattice structure
- resulting cohesive energy is the lattice energy

150
Q

crystal structures

A

-The crystal structure is determined by how many anions can fit around each
cation
-Ions with higher charges attract (or repel) more strongly than those with
lower charges
-Smaller ions attract (or repel) more strongly than larger ions because the
charges are closer to each other

151
Q

close-packed structures

A
  • hexagonal close packed (hcp)

- cubic close packed (ccp)

152
Q

interstices of close-packed structures

A

tetrahedral and

octahedral (giving a coordination for the cations of 4 or 6, respectively)

153
Q

NaCl structure

A 
aCl Structure
• Stoichiometry: MX
• 0.73 > r₊/r₋ > 0.41
• Coordination: 6:6
• Consists of octahedral arrangement of one ion around the other
154
Q

CsCl structure

A
  • Stoichiometry: MX
  • r₊/r₋ > 0.73
  • Coordination: 8:8
  • Consists of interpenetrating cubic arrangements of Cs+ and Cl
155
Q

ZnS - Wurtzite structure

A 
• Stoichiometry: MX
• 0.41 > r₊/r₋ > 0.22
• Coordination: 4:4
• Consists of tetrahedral arrangements of
one ion around another based on a hcp
arrangement of S²⁻ ions
156
Q

ZnS - Zinc blende structure

A 
• Stoichiometry: MX
• 0.41 > r₊/r₋ > 0.22
• Coordination: 4:4
• Consists of tetrahedral arrangements
of one ion around another based on a
ccp arrangement of S²⁻ ions
157
Q

lattice energy

A

the change in enthalpy* associated with

separating 1 mole of ionic solid into gaseous ions

158
Q

∆𝐻ₗₐₜₜᵢ𝒸ₑ (lattice energy)

A

∆𝐻ₗₐₜₜᵢ𝒸ₑ = 𝑧₊z₋𝑒²/4𝜋𝜀₀𝑟 x M

159
Q

lattice energy trends - moving down group

A
  • n gets bigger, so atom gets bigger
  • Ionic radius related to size of atom (for both cations and anions)
  • Electrostatic energy ∝ 1/r
  • Lattice energy decreases
160
Q

lattice energy trends - moving across period

A
  • Ionic charge changes
  • Singly charged (e.g. group 1A, 7A) or doubly charged (e.g. group 2A, 6A) ions
  • Lattice energy is ∝ (𝑧₊ x z₋)
  • Big increase in lattice energy when doubly charged ions involved
161
Q

consequences of lattice energy

A
  • ionic solids are hard, rigid, brittle
  • v high melt + boil points bc large energy req to move ions from fixed positions and separate them
  • ionic solids do not conduct electricity, as ions are fixed. Ions only free to move in molten or solution form
162
Q

Lewis symbols: dots

A

represent valence electrons

163
Q

pairs

A

bonding pairs and lone pairs

164
Q

Writing lewis structures

A

1- place atoms relative to each other

2- add up valence electrons

3- draw single bond frome each surrounding atom to central atom

4- distribute remaining pairs of electrons

165
Q

resonance structures

A

-For some molecules with double bonds next to single bonds we can write more than one
Lewis structure
-typically indicated with double headed arrow

166
Q

benzene: electron-pair delocalisation

A

electron density spread over a few atoms

167
Q

Bond order formula

A

Bond Order = no. of shared electron pairs/no. of bonds

168
Q

drawing molecules: solid wedge

A

bond/group projecting out towards viewer

169
Q

drawing molecules: hashed wedge

A

bond/group receding away from viewer

170
Q

shapes of molecules

A
  • linear
  • trigonal planar
  • tetrahedral
  • trigonal bipyramidal
  • octahedral
171
Q

linear shape

A

180°

172
Q

trigonal planar shape

A

120°

173
Q

tetrahedral shape

A

109.5°

174
Q

trigonal bipyramidal shape

A

90° and 120°

175
Q

octahedral shape

A

90°

176
Q

classifying molecular shapes using designation

A

AXₘEₙ

A = central atom
X = surrounding atom
E = nonbonding groups
n,m = integers
177
Q

simplest examples of each shape

A

Linear: AX₂

Trigonal Planar: AX₃

Tetrahedra: AX₄

Trigonal bipyramidal: AX₅

Octahedral: AX₆

178
Q

lone pairs and bonding pairs

A

In general lone pairs will repel bonding pairs more than bonding pairs repel each other
So, a lone pair will decrease angle between bonding pairs

179
Q

shapes incl non-linear shapes

A
  • linear
  • trigonal planar
  • tetrahedral
  • trigonal bipyramidal
  • octahedral
  • v-shaped/bent shape
  • trigonal pyramidal
180
Q

examples of breaking the octet rule

A
  • too many electrons
  • odd number of electrons
  • too few electrons
181
Q

breaking octet rule: too many electrons examples

A

eg. PF₅, SF₆, XeF₄

182
Q

breaking octet rule: too many electrons reason

A
  • electronegativity of F large enough, can form electron-pair bonds w/ central atom using electrons expected to be lone pairs if octet rule obeyed.
  • The energy change from satisfying octet rule for extra F outweighs effect of breaking it for central atom – can also occur for other very electronegative atoms (e.g. O, Cl) expanding the octet
183
Q

molecular shape - summary of common shapes

A

in slides 9

184
Q

odd number of electrons:

A

Does not have all its electrons in pairs.

Called free radicals + are extremely reactive

185
Q

Too few electrons: examples

A

eg. BeF₂, BF₃

186
Q

Too few electrons: reason

A

Be only has two valence electrons and so can only form two shared electron pair bonds – Similarly B can only form three shared electron pair bonds. As surrounding halogens are much more electronegative they will not form multiple bonds to central atoms

187
Q

too many electrons: where this expansion occurs

A

only occurs for nonmetals from period 3 or higher – requirement that d orbitals be available

188
Q

Differences between covalent and ionic substances

A
  • most covalent are comprised of molecules
  • ionic substances do not consist of molecules - continuous crystals in 3D
  • ionic bonding involved attraction between oppositely charged ions
  • covalent involves mutual attraction between + charged nuclei and - charged electrons - sharing
189
Q

covalent bonds

A

arise from sharing of pairs of electrons between atoms, to form molecules

190
Q

bond energy

A

defined as the standard enthalpy* change (ΔH°) for breaking the bond in 1 mol of gaseous molecules

191
Q

breaking bonds - endo or exo thermic?

A

an endothermic process, so energy always positive

192
Q

ΔHᵣₓₙ

-standard enthalpy of a reaction

A

ΔHᵣₓₙ = Σ(all bond energies in reagents) – Σ(all bond energies in products)

dont forget to factor in moles when you do these calculations

193
Q

haber process

A

-formation of ammonia:

N₂ + 3H₂ –> 2NH₃

-exothermic reaction

194
Q

bond length

A

of a covalent bond is the distance between the nuclei of two bonded atoms

195
Q

bond lengths and atomic radii

A

-Bond lengths are related to the atomic radii of the atoms bonding

196
Q

bond lengths and interdependence

A

Higher bond order results in shorter bond length (and higher bond energy)

eg. C-C triple bond is shorter than C-C double or single bond

197
Q

large bond energies

A

suggest very strong bonds

198
Q

Gases, liquids, low melting temp solids and bonds

A
  • Strong bonding forces within a molecule

- But weak intermolecular forces between molecules

199
Q

boiling a liquid

A

by boiling a liquid molecule we are overcoming these relatively weak intermolecular forces
(rather than the strong covalent bonds)

200
Q

seven pure elements that form diatomic molecules

A
  • Hydrogen (H₂)
  • Nitrogen (N₂)
  • Oxygen (O₂)
  • Fluorine (F₂)
  • Chlorine (Cl₂)
  • Iodine (I₂)
  • Bromine (Br₂)
201
Q

diatomic atoms

A

consist of 2 atoms bonded together

202
Q

valence bond theory

A

Covalent bond forms when orbitals of two atoms
overlap and a pair of electrons occupy the overlap region

Molecular shapes rationalised through the formation of hybrid orbitals

203
Q

molecular orbital theory

A

Involves molecular energy levels

A quantum mechanical treatment of molecules involving molecular orbitals

204
Q

two main ways for orbitals to overlap

A

sigma bond

pi bond

205
Q

sigma bond

A

end to end

206
Q

pi bond

A

side to side

207
Q

Molecular Orbital Theory Bond Order Formula

A

𝐵𝑜𝑛𝑑 𝑜𝑟𝑑𝑒𝑟 = 1/2 [(𝑛𝑜. 𝑜𝑓 𝑒⁻ 𝑖𝑛 𝑏𝑜𝑛𝑑𝑖𝑛𝑔 𝑀𝑂) − (𝑛𝑜. 𝑜𝑓 𝑒⁻ 𝑖𝑛 𝑎𝑛𝑡𝑖𝑏𝑜𝑛𝑑𝑖𝑛𝑔 𝑀0)]

208
Q

single bonds

A

observed when the no. of electrons in bonding orbitals is two (or one pair) greater than the number in antibonding orbitals (i.e. one shared pair)

209
Q

double bonds

A

observed when four electrons more (or two pairs) are in bonding orbitals

210
Q

triple bonds

A

observed when we have six more electrons (or three pairs) in bonding
orbitals

211
Q

bond order > 0

A

molecule more stable than separate atoms i.e. it will form

212
Q

higher bond order

A

higher bond order = stronger bond

213
Q

polar covalent bond

A

where the more electronegative atom acquires a partial negative charge (δ-) and the less electronegative atom acquires a partial positive charge (δ+)

214
Q

Polar bond

A

has partial ionic character – in fact, ionic bonds can be considered as an extreme case of polar bonding!

215
Q

electronegativity difference (Δχ)

A

difference in electronegativity between the bonded atoms is directly related to the bonds polarity – a higher Δχ indicates a higher degree of ionic character

216
Q

dipole moment

A

-dipole moment associated with a polar bond

μᵦ = δL

where L is bond length

217
Q

dipoles measurement

A

measured in Debye (D)

where 4.8D is dipole arising from charges of e⁺ and e⁻ separated by 0.1nm

218
Q

molecular dipole moment formula

A

μ = Σ(μᵦ) + Σ(μₗₒₙₑ ₚₐᵢᵣₛ)

219
Q

polar molecule

A

one with a non-zero dipole moment (μᵦ) – this arises from the polarity of the bonds in the molecule and the presence
of directional lone pairs

220
Q

ability of a solvent to dissolve ionic/polar substances

A

depends on its

molecular dipole moment

221
Q

other properties that depends on molecular dipole moment

A

dielectric constant

refractive index

222
Q

polar bonds and polar molecules

A

polar bonds will not always result in a polar molecule

revise this

223
Q

polar molecule

A

one with a permanent dipole moment

224
Q

interaction of water with an ionic substance eg. NaCl

A

attraction between the anion (Cl-) and the H end of the water molecule

attraction between the cation (Na+) and the O end of the water molecule

225
Q

main reason ions become separated in solution

A

the ion-pole attractions are greater than the ion-ion attractions

226
Q

electron sea model

A
  • explains properties of metals
  • Atoms contribute their valence electrons to a delocalised electron sea – resultant cations arranged in an ordered array
  • Held together by attraction of cations for this electron sea
227
Q

how electron sea is diff from ionic and covalent bonding

A
  • no anions so metal ions are not locked in place like in ionic crystals
  • No particular pair of atoms bonded through shared localised electron pair – all cations attracted to delocalised electron sea
228
Q

metals: going down a group

A

larger ions – weaker attraction to electron sea

229
Q

metals: going across a period

A

more valence electrons – greater attractive forces between cations and electron sea

230
Q

explaining properties of metals: conductivity

A

abundance of mobile electrons (negative charge carriers) makes metals good electrical conductors

231
Q

explaining properties of metals: mechanical properties

A

metal cations are not held rigidly.

If deformed, can slide over each other + adopt new positions
– so metals can dent and bend.
-Softer metals are ductile and malleable

232
Q

explaining properties of metals: melting and boiling points

A

-moderate melting points and high boiling points
– in liquid form metals’ cations still attracted to electron sea
– in gaseous form cation + associated electrons must be liberated w/ a high associated energy cost, so high boiling points

233
Q

alloying

A

involves mixing a molten metal with another molten metal

234
Q

types of alloys

A

substitutional alloys

interstitial alloys

235
Q

substitutional alloys

A

form when the atoms are similar in size – effectively substituting one metal for another

236
Q

interstitial alloys

A

formed when one of the atoms involved is considerably

smaller than the other (typically a non metal) and sits in interstitial positions between the “solvent” metal atoms

237
Q

effect of temp on conductivity: metals

A

electrical conductivity is high + decreases as temperature (T) increases.
• main source of resistance is vibration of atoms in the solid which scatter electrons – as T goes up, amplitude of these vibrations increases.

238
Q

effect of temp on conductivity: semiconductors

A

electrical conductivity is relatively low + increases as T increases.
• Raising temp creates thermally excited electrons which can “jump over” band gap.

239
Q

how impurities affect conductivity

A
  • In metals impurities will scatter electrons + decrease conductivity
  • in semiconductors, impurities can increase conductivity through process known as doping
240
Q

doping

A
  • silicon most common semiconducting material
  • conductivity can be improved through doping
  • doping leads to intro of states in band gap
241
Q

n-type doping

A

When Si doped with P or As or another 5A element

242
Q

p-type doping

A

When Si doped with B or Ga or another 3A element

Ch (3 decks)

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