Chemistry Unit 2

1
ChemistryUnit 2

• Chapters 4-6
• Atomic Properties and Bonding

2
First Things First

• Lab reports
• Give back homework
• Tests

3
The Modern Atom

• Bohr Model
• Atoms made of 3 particles

4
Light

• Originally thought to be a wave
• Discovered to also behave like a particle
• A packet of energy called a quantum is wavelike
• A photon contains a quantum and behaves like a
  particle

5
Wave-Particle Duality

• Developed by Einstein
• Light behaves both as a wave and a particle
• Quanta act like waves and are a unit of energy
• Photons act as particles and are made of quanta
• Photons have 0 mass and are pure energy

6
Electromagnetic Radiation

• Energy that acts as a wave
• Energy linked to wavelength and frequency of the
  light

7
Electromagnetic Radiation

• All waves move at 3.00 x 108 m/s in a vacuum
• Wavelength (?) is the distance between
  corresponding points on a wave
• Frequency (?) is the number of waves per second
  that pass by a point

8
Energy of Waves

• High frequency and low wavelength are high energy
• Low frequency and high wavelength are low energy

9
Electromagnetic Spectrum
10
Electromagnetic Spectrum
11
Visible Light

• Only certain range detectible to human eyes
• 400-700 nm or 5 x 1014 1 x 1015 Hz
• Violet is the highest energy, red the lowest

12
Ultraviolet and Infrared

• Ultraviolet (UV) means more than violet, higher
  energy and frequency, shorter wavelength
• This is why UV rays are damaging
• Infrared (IR) is less than red, lower energy and
  frequency, longer wavelength
• This is why IR remotes dont do damage

13
Electromagnetic Spectrum

• Be able to rank from high to low by wavelength,
  frequency, or energy
• Pg 98

14
Relations of Wavelength and Frequency

• The speed of light (m/s) is equal to the
  wavelength (m) times the frequency (1/s)
• c ??
• Which is where a physics joke comes from.

15
The Photoelectric Effect

• Explains part of the particle nature of light
• When light shines on some metals they emit
  electrons

16
The Energy of Light

• Quantum also happens to be energy gained or lost
  by electrons as they move closer or further from
  the nucleus, emitted as photons
• Energy (J) equals Plancks Constant (Js) times
  the frequency (1/s)
• E h?

17
Plancks Constant

• Developed by Max Planck
• Relates the energy to the frequency
• 6.626 x 10-34 Js

18
Line-Emission Spectra

• Energy added to gas atoms and electrons emit
  light waves
• Electrons promoted from the ground state to an
  excited state
• Ground state is the lowest energy level
• Excited state is any higher energy state

19
Line-Emission Spectra

• Each element has a unique set of lines

20
Line-Emission Spectra

• Lyman Series
• UV
• Balmer Series
• Visible, most commonly used
• Paschen Series
• IR

21
Quantum Theory

• When electrons lose energy they emit it as
  photons
• Studies the behavior of atoms and their energies
  in relation to photons and quanta of energy

22
Niels Bohr

• Proposed set orbitals for electrons
• These were based on the line-emission spectrum of
  different elements
• Photons emitted have specific energies based on
  how much energy is lost in falling to a lower
  energy level

23
Energy Levels

• Electrons can gain or lose energy and move up or
  down in energy states
• Each element has its own set of allowed energy
  levels
• Biggest difference in energy levels Lyman
• Moderate difference in energy levels Balmer
• Small difference in energy levels Paschen

24
Energy Levels
25
Energy Levels
26
Wave-Particle Duality (Again)

• 1924 Louis de Broglie realized electrons also
  acted like a wave as well as a particle
• Waves confined to a set space have only certain
  frequencies, much like electrons do
• de Broglie pointed out that electrons acted like
  waves confined to the space around an atom

27
de Broglie Waves

• All matter can behave as a wave
• Wavelength of a particle determined by the mass
  and speed of the particle
• ? h/mv

28
Detecting de Broglie Waves

• Electrons detected by photons, but act similarly
  to photons (wave-particle duality)
• Any attempt to find where an electron is changes
  where it is
• A solution proposed by Werner Heisenberg

29
Heisenberg Uncertainty Principle

• It is impossible to determine both the position
  and velocity of a particle.
• Also stated Observing a system disturbs it in
  such a way that any data collected cannot be
  certain.

30
Schrödingers Equation

• Erwin Schrödinger developed an equation to
  mathematically show electrons act as wave
• Solutions prove that only certain energies are
  allowed
• Gives probability regions, called orbitals, where
  electrons can be
• Probability regions are where electrons are 95
  likely to be

31
Solutions to Schrödingers Equation

• Gives 4 quantum numbers
• Principal quantum number n
• Angular momentum quantum number l
• Magnetic quantum number m
• Spin quantum number

32
Principal Quantum Number

• n gt 0
• Gives the energy level of the electron
• As n increases so does energy and distance from
  the nucleus
• Electrons allowed per level is 2n2 up to 32

33
Principle Quantum Number
34
Angular Momentum Quantum Number

• 0 l n-1 3
• Gives the sublevel
• All sublevels but n1 have multiple orbitals

35
Angular Momentum Quantum Number
36
s Sublevel Orbitals

• Spherical
• Holds 2 electrons

37
p Sublevel Orbitals

• 3 orbitals
• Named by which axis they lay along
• Holds 6 electrons total

38
d Sublevel Orbitals

• 5 orbitals
• Named for orientation
• Holds 10 electrons total

39
f Sublevel Orbitals

• 7 orbitals
• Dont worry about what they look like or are
  named
• Holds 14 electrons total

40
f Sublevel Orbitals
41
Orbitals

• s is lowest energy, then p, d, and f

42
Magnetic Quantum Number

• -l m l
• Gives orientation of the orbitals
• Gives number of allowed orbitals
• 2l 1

43
Spin Quantum Number

• Only 2 possible values, relating to the two
  magnetic poles
• 1/2 or -1/2
• Each orbital can have 2 electrons and they must
  have opposite spins

44
Game

• Pick a partner and sit together
• You can use your book and planner, just not other
  teams papers
• If you dont have a calculator get one from the
  bench
• The team with the highest ratio of right/time
  gets an ECP

45
Game

• Go over answers

46
Electron Configurations
47
Electron Configurations

• Shows how electrons are arranged
• Gives energy level, sublevel, and how many
  electrons

48
3 Rules to Placing Electrons

• Aufbau Principle
• Pauli Exclusion Principle
• Hunds Rule

49
Aufbau Principle

• Naturally, electrons are found in the lowest
  energy levels possible

50
Aufbau Principle
51
Pauli Exclusion Principle

• Each electron has to have a unique set of quantum
  numbers
• AKA Electrons must be paired by opposite spins.

52
Pauli Exclusion Principle

• 1/2 and -1/2 spin are paired
• Each orbital holds 2 electrons

53
Pauli Exclusion Principle
54
Hunds Rule

• Electrons being placed in sublevel orbitals of
  equal energy will be placed unpaired first.
• All unpaired electrons must have the same spin.

55
Hunds Rule
56
Practice

• Following Hunds Rule, how many unpaired
  electrons are allowed in each sublevel?

57
Orbital Notation

• Shows electrons as up or down arrows,
  representing 1/2 and -1/2 spin states
• Orbitals as lines (sometimes boxes) labeled with
  energy level and sublevel

58
Practice

• Using the diagram, and what you know about
  orbitals, and show atoms with 3, 7, and 18
  electrons

59
Electron Configuration Notation

• Uses energy level diagram and orbital notation to
  write out the electron configuration in the form
  nlx
• n is principle quantum number (energy level), l
  is the sublevel (letter) derived from the angular
  momentum number, and x is the number of electrons
  in that sublevel

60
Refresher

• How many electrons are allowed in each sublevel?
• s
• p
• d
• f

61
Example

• H is 1s1
• He is 1s2
• C is 1s2 2s2 2p2
• Kr is 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6

62
Why does the d Block Start at 3?

• 3d is higher energy than 4s, but less than 4p
• s and p are always the same as the period, d is
  one less, f is 2 less

63
Writing Electron Configuration Notation

• Work up from atomic 1
• Electrons shown in notation always equals the
  atomic number of the element for neutral atoms
• Use the periodic table as your guide

64
Writing Electron Configuration Notation

• Start at 1 and move across a period

65
Noble Gas Notation

• Written as previous noble gas in brackets and
  then the electron configuration not included in
  the noble gas
• C is He 2s2 2p2

66
Practice

• Write the following in noble gas notation
• Germanium
• Strontium
• Tin

67
The Periodic Table
68
The Periodic Table

• Periodic means repeating in a pattern
• The Periodic Table of the Elements arranges
  elements by increasing mass and of protons

69
The Periodic Table

• 2 groupings
• Groups Vertical columns, elements in groups
  have similar properties
• Periods Horizontal rows, show similar electron
  configurations and energy levels

70
Famous Dead Guys

• Dmitri Mendeleev
• Russian 1834-1907
• Arranged the 63 known elements by mass on a
  periodic table and in groups by properties

71
Mendeleev

• Based on properties of the groups he could
  predict the masses and properties of undiscovered
  elements
• Arranged in order of mass, and noticed a pattern.

72
The Periodic Table
73
Group 1

• Alkali Metals
• 1 electron in the s sublevel
• Malleable, ductile, good conductors, very soft
  metals, very reactive, explosive in water, cesium
  and francium are some of the most reactive
  elements, never found free in nature

74
Group 2

• Alkaline Earth Metals
• 2 electrons in the s sublevel
• Malleable, ductile, good conductors, very
  reactive, but less than alkali metals, never
  found free in nature

75
Groups 3-12

• Transition Metals
• Central block of metals
• 2 electrons in the s sublevel and 1-10 in the d
  sublevel
• Malleable, ductile, good conductors, some can
  produce magnetic fields, have unusual electron
  shells

76
Group 13-15 Metals

• Very dense solids
• Bottom left corner of the p block
• Only 7
• Aluminum, Gallium, Indium, Tin, Thallium, Lead,
  and Bismuth

77
Group 13-16 Metalloids

• Found on the line between metals and nonmetals
• What is a metalloid?
• Only 7
• Boron, Silicon, Germanium, Arsenic, Antimony,
  Tellurium, Polonium

78
Group 14-16 Nonmetals

• Upper right of 14-16
• Nonmetals
• Only 6
• Carbon, Nitrogen, Oxygen, Phosphorus, Sulfur, and
  Selenium

79
Group 17

• Halogens
• Only group to have solids, liquids, and gases
• Have 2 electrons in the s sublevel, some have d
  sublevel, all have 5 electrons in the p sublevel
• Nonmetals, insulators, form salts

80
Group 18

• Noble Gases
• Inert (dont react)
• The most stable elements
• Full outer electron shell (valence shell)

81
Lanthanides and Actinides

• 2 periods at the bottom of the chart
• Also called Rare Earth Metals
• Have the f sublevel
• Most are synthetic
• Most are radioactive
• Very heavy and unstable

82
The Modern Periodic Table

• There were a few places where Mendeleevs table
  didnt quite work
• This was fixed by Henry Moseley in 1911

83
Famous Dead Guy

• Henry Moseley
• English 1911
• Working in Rutherfords lab
• Developed the Periodic Law
• Studied spectra of elements
• Found that the elements should be arranged by
  protons, not mass

84
Periodic Law

• The physical and chemical properties of the
  elements are functions of their protons and
  electrons, not mass.

85
Periodicity and Electron Configurations
86
Periodicity and Electron Configurations

• According to the Periodic Law the properties and
  periodicity are a function of the protons and
  electrons of an atom

87
Atom Radii

• Defined by ½ the distance between the nuclei of
  identical atoms bonded together
• Radii get smaller across a period and larger down
  a group
• Smallest radii are at the upper right, largest at
  the lower left

88
Atomic Radii
89
Practice

• Which in each group has the larger radius
• Ba or Mn
• In or Sn
• O or S

90
Ionization Energy

• The energy required to remove an electron from an
  atom
• The stronger the electron is held to the atom the
  higher the ionization energy
• Only deals with loss of an electron
• Increases across the period and decrease down the
  group

91
Electronegativity

• Developed by Linus Pauling
• Measures the ability for an atom in a compound to
  attract electrons from another atom in the
  compound
• AKA Measures how badly they want another electron

92
Electronegativity

• High electronegativities mean they will more
  readily take an electron, lower means less likely
• Electronegativities near or below 2 means they
  will readily give up electrons

93
Electronegativity

• Tend to increase across a period and decrease
  down a group
• Upper right of the p block will very easily take
  electrons
• Lower left of s block easily lose electrons

94
Valence Electrons

• Electrons in the outer electron shell
• Highest energy level electrons
• Interact in chemical reactions and compounds
• Up to 8 in main block elements
• The s and p blocks
• Called an octet
• Group patterns emerge

95
Practice

• How many valence electrons?
• H
• Ca
• Al
• Te
• Kr
• Mo

96
d and f Block Properties
97
d and f Block Properties

• All transition metals have similar properties
• Almost like theyre all part of the same group
• Can lose or gain several different numbers of
  electrons

98
d and f Block Properties

• Similar properties because they all have the same
  number of valence electrons
• Remember d block energy level is 1 less than
  period

99
d and f Block Properties

• The d block
• Radii decrease across a period, but less than
  main block
• The f block
• Radii fluctuate across a period
• Caused by filling orbital 2 energy levels less
  than valence

100
d and f Block Properties

• Ionization energies increase across period
• Ionization energies increase down the groups,
  unlike main block elements

101
d and f Block Properties

• Electronegativities are weak, prefer to give up
  electrons
• Follow same trends as main block elements

102
Ions
103
Ions

• Charged atoms
• Caused by adding or removing electrons
• Cation positive ion
• Lost electron(s)
• Anion negative ion
• Gained electron(s)

104
Ions

• Atoms gain or lose electrons to be more like
  noble gases
• Will usually only lose valence electrons
• Charge equal to number of electrons gained or lost

105
Practice

• Will these form cations or anions?
• Li
• Ba
• Sb
• F
• W

106
Practice

• What is the charge on these ions when they are
  like noble gases?
• Mg
• O
• P
• Fr

107
Ionic Radii

• Cations smaller than their neutral version
• More protons than electrons, held tighter
• Anions larger than their neutral version
• More electrons than protons, held less tightly
• Increase down a group

108
Ionic Radii
109
Practice

• What is the smallest ion?
• Are these ions smaller or larger than their atom?
• Li
• O2-
• I-
• Pd2

110
Chapter 5 Review Worksheet
111
Bonding
112
Bonding

• The periodic table is like a parts list or a
  buffet
• Atoms mix together to make compounds
• Chemical bonds are the attraction of atoms and
  interactions of valence electrons

113
Why bonding?

• Atoms are most stable when they have full octets
• Either share or take electrons

114
Ionic Bonding

• One atom steals the electrons of another
• Attraction between anions and cations
• Large difference in electronegativies
• ?EN gt 1.7
• Usually s block and nonmetal

115
Ionic Bonding

• Na has EN of 0.9
• Cl has EN of 3.0
• ?EN of 2.1
• ?EN ENanion Encation
• Use table on page 161

116
Covalent Bonding

• Sharing electrons
• Small difference in EN
• ?EN lt 1.2

117
Covalent Bonding

• Cl2
• ?EN 0
• HCl
• H has EN of 2.1
• Cl has EN of 3.0
• ?EN of 0.9

118
Polar Covalent Bonding

• Unequal sharing
• Moderate difference between atoms
• 1.2 lt ?EN lt 1.7

119
Polar Covalent Bonding

• LiI
• Li has EN of 1.0
• I has EN of 2.5
• ?EN of 1.5

120
Metallic Bonding

• Lots of cations (metals that have lost electrons)
  all have electrons flying around between them
• Occurs only between 2 or more metals
• These arrangements make metals malleable,
  ductile, and good conductors
• Malleable ability to be hammered into sheets
• Ductile ability to be drawn into wires

121
Bonding Visuals
122
Bonding Visuals
123
Bonding Visuals
124
Practice

• What type of bonds are these?
• NaAt
• FrBr
• MgCl2
• AgBr2
• HgO
• ZnCu

125
Displaying Bonds
126
The Octet Rule

• Atoms are most stable when they have a full octet
• Outer electron shell (energy level) is filled

127
Practice

• Draw the orbital notation of these atoms and show
  that they are more stable with overlapping
  orbitals
• HCl
• H2O
• NaCl
• MgF2

128
Exceptions to the Octet Rule

• Hydrogen and helium only need 2 valence electrons
• Boron likes 6
• Expanded valence atoms can have more than 8
  valence electrons when bonded with very
  electronegative elements

129
Electron Dot Diagrams

• Shows valence electrons as dots surrounding the
  elements symbol
• Use 2 dots per side, for up to 8 total

130
Practice

• H
• Sr
• Bi
• O
• Ar
• B

131
Lewis Structures

• Show molecules as electron dot structures with
  shared electrons shown between them
• Unshared electrons are called lone pairs
• 2 shared electrons can be shown with a dash
• 1 dash (2 electrons) is a single bond
• 2 dashes (4 electrons) is a double bond
• 3 dashes (6 electrons) is a triple bond

132
Drawing Lewis Structures

• Find how many electrons each atom wants to gain
  or lose
• Draw each atom separately
• Combine them and show shared electrons as dashes

133
Example
134
Practice

• Draw the following as Lewis Structures
• O2
• CH4
• C2H5OH
• N2

135
Resonance Structures

• Compounds with multiple Lewis Structures
• Reality is a combination of structures
• Resonance adds to stability

136
Example

• Ozone
• O3

137
Practice

• C4H6
• C5H8
• C6H6

138
Structural Formula

• Doesnt show dots for lone pairs, only dashes for
  bonds
• Carbons with all 8 electrons shared dont have to
  be shown, just as points

139
Practice

• HCl
• H2O
• CO2
• C6H6

140
Ionic Bonding
141
Ionic Bonding

• Ionic bonds are bonds where the difference in
  electronegativities is great enough that one atom
  steals the electron from another

142
Crystals

• Many ionic compounds for crystals
• A 3D structure of cations and anions held together

143
Formula Units

• Shows the simplest ratio of cations and anions
• Since they form crystals the real structure is
  usually a multiple of this
• NaCl is a 11 ratio, but is found in large
  crystals

144
Why Crystals

• The regular arrangement of the ions in a crystal
  have the lowest energy
• This makes them more stable
• The structure of the crystal (where the ions fit
  in) is called a crystal lattice

145
Lattice Energy

• The energy released when ions as a gas form an
  ordered crystal
• Negative energy is a release of energy, this is
  favorable
• The more negative the value the more stable the
  crystal

146
Forming Ionic Bonds

• Ionic compounds are shown as ions in Lewis
  Structures
• Final structures must be stable

147
Example
148
Ionic Properties
149
Ionic Properties

• Atoms held together by attraction of positive and
  negative charges
• Molecules barely held together

150
Ionic Properties

• Hardness of the compound, melting point, and
  boiling point all depend on how well the formula
  units are held together
• Stronger attractions harder, higher bp and mp

151
Ionic Properties

• Most are solids at room temperature
• Strong attraction between formula units
• In solutions (liquids, usually water) they
  dissociate
• Dissociate means falls apart
• Cations and anions found seperately

152
Visual
153
Polyatomic Ions
154
Polyatomic Ions

• Cations or anions made of multiple atoms
• A group of covalently bonded atoms with a net
  charge

155
Polyatomic Ions
156
3D Chemistry
157
VSEPR Theory

• Valence Shell Electron-Pair Repulsion Theory
• The repulsion of the electrons to each other
  makes the bonds between atoms desirable to be as
  far apart as possible.

158
VSEPR Theory

• Lone pairs need to be accounted for as well
• Geometry can be determined based on the number of
  atoms and lone pairs bonded to the central atom
  (stearic ) and the number of lone pairs

159
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160
Example

• Draw the Lewis Structure for water (H2O)
• According to VSEPR Theory how do we get the
  electrons farthest apart?
• Hint Think 3-Dimensionally
• 4 bonds, 2 lone pairs
• Bent Structure

161
Practice

• NH3
• CO2
• NaAt
• PCl3
• HCP
• CH4

162
Bond Hybridization
163
Hybridization

• Geometry determined by hybridization of the bond
• A hybrid bond is a combination of orbitals to
  create a new orbital

164
Example

• Draw methanes (CH4) orbital notation
• 2s and 2p merge to form a 2sp3
• An sp3 sublevel has 4 orbitals, as many as there
  were before, just at a single energy

165
Hybridization
166
Intermolecular Forces
167
Intermolecular Forces

• Forces between compounds, not within
• Forces within a compound are intramolecular

168
Dipole-Dipole

• Strongest force between polar molecules
• Uneven distribution of charges
• Polar covalent bonding

169
Hydrogen Bonding

• A specific type of dipole-dipole force
• Only occurs in H-O, H-F, and H-N bonds
• Hydrogens are attracted to the lone pairs of
  another molecule

170
London Dispersion Force

• Attraction due to momentary dipole created by the
  natural movements of electrons
• Very weak, occurs in all atoms and compounds

171
End of Unit 1

• Questions?