Gr. 10 Andries Oliviers' CAPS aligned Physical Sciences Contains: States of matter Periodic table Chemical bonding

1. Matter and materials

2. Properties of matter
Macroscopic - Microscopic – Not
Observed with observed with
senses senses
Strength Atom that bonds
Thermal and electrical
conductivity consists of
Magnetic properties Kinds of bonds
Solubility Forces between
Malleability and ductility bonds
Density
MP and BP

3. Mixtures

4. Properties of mixtures
 The substances keep their original properties.

 The substances don’t have to be mixed in a
fixed ratio.

 The mixture can be separated with simple
methods.

5. Two mixtures
Homogenous Heterogenous
There are more than one
phase.
The substances are all Eg sand and water
in one phase Suspension – solids that
float in liquids.
Eg air
Eg muddy water
You cannot identify the Emulsion – steady mix of
different components insoluble substances in a
of the mixture. liquid.
Eg milk, mayonnaise

6. Separation methods
Because mixtures retain the original
properties, we can separate the different
substances by physical methods

7. Homogenous separation
Evaporation
Distillation
Fractional distillation
Chromatography
Centrifugation

8. Evaporation

9. Chromatography
Taken from regentsprep.org

10. Centrifuge
Taken from faqs.org

11. Heterogenous separation
Filtration
Separating funnel
Decanting
Sorting
Sieving

12. Filtration

13. Sorting
Taken from codeunit.co.za

14. Sieving
Taken from sbs.utexas.edu

15. Pure substance
Pure substances are made of only one
substance or a compound.

16. Pure substances
Elements Compounds
A compound consist of two or
An element only consists more different elements
of the same atoms. bonded together.
An element can’t be Can only be separated by
separated into simpler chemical methods.
substances. Compounds’ properties differ
Elements are categorised from the individual elements.
as metals, non-metals and Joined in fixed ratios.
semi-metals. Chemical reactions occur
during formation.

17. Periodic table

18. Formulae
H2O is the formula of water
O

H 2O
H H
2 hydrogen atoms
1 oxygen atom
 CaO, CaSO4, Ca(OH)2, NH4NO3, CO2, NH3
 Diatomic molecules
 Ionic bonds

19. Ionic bonds
Group 1, 2, 6, 7
Transition metals – give off 1, 2, or 3 electrons
Mono-atomic anion – ‘ide’
Polyatomic ion/radical

20. Physical differences
Metals Non-metals
Metallic lustre Dull (except graphite &
Electrical conductors diamonds)
Thermal conductors Poor electrical and
Opaque thermal conductors
Malleable and ductile Some solids = opaque,
Solids @ room gases are translucent
temperature, except Hg Solids = brittle
High MP and BP Low MP and BP

21. Semi-metals
Generally have
properties of metals,
but a few non-metal
properties as well.
Ability to conduct
electricity increases
with heat. (in contrast
with metals)

22. Properties of semi-metals
Shiny or dull
Conduct heat and electricity better than non-
metals, but weaker that metals
When heated, they can conduct electricity
better.

23. Electrical conductors, Semi-
conductors and insulators
Metals – conductors
Non-metals – insulators
Semi-metals – poor conductors, called semi-
conductors

24. Thermal conductors or
insulators
Thermal conduction is the flow of heat from a
high temperature to a low temperature.
Metals – thermal conductors
Non-metals – thermal insulators. All materials
that trap air = poor conductors, because air is a
poor conductor.

25. Magnetic and non-magnetic
materials
Ferromagnetic
elements – strongly
attracted to magnets
Fe, Ni, Co, Alnico,
ceramic (insulating
magnets), magnetite

26. Phases of matter
And the kinetic molecular theory

27. The three phases of matter

28. SOLIDS
Solids keep its shape and can only be dented, broken or
bent. Hard
High density No compressibility
Fixed volume
Made up of small particles
Vibrate only
Very small spaces between particles
Strong attractive forces – causes specific shape
No diffusion
Have crystalline structure
Have a specific melting point

29. Melting
Solidification
Liquid

30. LIQUIDS
No fixed form Not hard
High density No compressibility
Flows Fixed volume
Particles move in ordered fashion
Collisions occur
Diffusion occur
Smaller spaces between particles than with gases
Exerts pressure in all directions
Weak force between particles
Specific freezing point and boiling point

31. Evaporation
Condensation
Gas

32. GASES
No fixed form Not hard
Low density Easily compressible
Flows No fixed volume
Particles move fast
Greater collision
Big open spaces
Weak/no forces between particles
Involuntary motion
Diffusion occurs
Exerts pressure in all directions

33. Kinetic model of matter

34. Motion of particles
Diffusion
… is the movement
of particles from a
high concentration to
a low concentration.

35. Motion of particles

36. Kinetic model of matter
1. All matter consist of small particles
2. Particles are in constant motion
3. Spaces between the particles
4. Constant collisions between particles and container
5. Temperature is a measure of the kinetic energy of
the particles
6. Forces between particles
7. Phase changes occur when energy changes occur

37. Phase changes
1. Condensation 2. Solidification
4. Evaporation 3. Melting

38. 1. Condensation
Prior to condensation:
* particles slow down
* not far apart
* less violent collisions
Phase change follows:
* Spaces decrease
* Forces increase
* more orderly arrangement

39. 2. Solidification
Prior to solidification:
* particles move very slowly
* particles very close to each other
* only vibrates
Phase change follows:
* very small spaces between particles
* forces between particles become very
strong
* orderly arrangement

40. 3. Melting
Prior to melting:
* particles move fast
* particles further apart
Phase change follows:
* spaces increase
* forces decrease
* less orderly arrangement

41. 4. Evaporation
Prior to evaporation
* particles move very fast
* particles very far apart
* violent collisions due to high speed
Phase change follows:
* spaces between particles are big
* forces negligible
* disorderly arrangement

42. Evaporation vs. boiling
BOILING EVAPORATION
Occurs @ B.P. Occurs @ temp below
Occurs throughout the B.P.
liquid Occurs only on surface
Quicker Slow
Temp remains constant Causes cooling → heat
during boiling absorbed from
environment

43. Heating curve of water

44. Heating curve of water

45. Atomic structure
Atomic models

46. The electrical nature of matter
Michael Faraday
Electrical current through salt solutions
Amount of Q = amount of atoms reacting

47. Dalton's atomic theory
Michael Faraday
Amount of Q = amount of atoms reacting
●

48. Thomson’s atomic model
Charge and mass of
electrons of all -
substances are the
same. -
Thus, electrons in all
substances must be the
-
same.
Substances differ -
because electrons are
arranged differently.

49. Rutherford’s atomic model
Bombard thin gold foil with α-particles (heavyweight kind of radioactivity).
Fluorecent zinc sulphide screen opposite foil = α-
particle detector.
fluorecent zinc sulphide screen
α-particles
gold foil
Radioactive
source

50. Rutherford continued
Observations
α – (+2)
 Most of the α-particles
showed no diversion. +
Atom
+ +
Some were deflected
 Some were reflected
+ + + +
Conclusions
 Electrons occupy the + + +
greatest volume of the
atom. + + + +
 Positive particles are
grouped together in the
nucleus – very heavy, but
small.

51. Assumptions
1. The positive charges are all together in a small
volume in the nucleus.
2. The nucleus is surrounded by a space that
contains the e- (v. Small mass) – e- are
responsible for the great volume of an atom.
3. Mass is concentrated in the nucleus.
Later investigations predicted that the nucleus is
positively charged. # Protons = # electrons. e-
don’t move like bees around a hive, e- would collapse
into nucleus.

52. Bohr’s atomic model
Electrons move in orbits
Electrons with the same energy move around
in the same orbit
Electrons in orbits further away from nucleus
have a higher energy

53. Planetary atomic model
e- move in energy levels
e- with same E values, move in
same E levels
Valence orbitals have higher
energy than those close to the
nucleus
Energy levels closer to the
nucleus are filled first with e-
Each energy level can only
take a specific amount of e-
e- in orbits close to the nucleus
are lower in energy level than • When e- absorb energy
orbits further away it rises in energy level.
If e- are in lowest possible • This (excited) state is
energy level – ground state unstable and e- fall back
to lower energy levels

54. Line spectra
Electrons in the ground state absorb energy. Electrons are
now excited, and move to a higher energy level.
This electron is now unstable.
It falls back to its ground state, radiating extra energy as
light.
The separate coloured lines of light show electrons only
have certain energy. Electron’s energy is thus quantised.
Each element has it’s own unique line spectrum.
Line spectra occur when gases are heated of an electric
current is passed through it.

55. Wave mechanical atomic model
Bohr’s atomic model explains the structure of
hydrogen, but not those of atoms with more than
one electron.
The discovery of wave properties of electrons
gave us a more acceptable model.
e- have both particle and wave properties.
Schrodinger stated that moving e- form a 3D wave
space that surrounds the nucleus, called an orbital.

56. Neutron
J. Chadwick discovered a particle with a mass
nearly equal to the proton.
Neutral charges, called neutrons.

57. Atomic mass and diameter
Atoms are extremely small with small masses.
ELEMENT AVG ATOMIC MASS
Hydrogen 1,673 55 X 10 -27
Carbon 1,994 36 X 10 -26
Oxygen 2,656 59 X 10 -26
Uranium 3,952 33 X 10 -25
Diameters are also extremely small.
Most of the volume of an atom is empty space, the
nucleus accounts for most of the mass of an atom at
the centre.

58. Relative atomic mass
Hydrogen is the lightest atom and is chosen as the standard for an
atomic mass scale.
This mass is equal to 1.
Using proportion to find the atomic masses of other elements
relative to a mass of 1.
1,673 55 X 10 -27 kg of Hydrogen = 1 on the Hydrogen scale
So, 2,656 59 X 10 -26 of oxygen = 1x2,656 59 X 10 -26 kg
1,673 55 X 10 -27 kg
= 15, 87 on the Hydrogen
scale

59. Structure of the atom
The atom consists of very many small subatomic particles.
In chemistry we work with protons, neutrons and electrons.
Protons and neutrons are in the nucleus at the centre.
Electrons occupy a large region around the nucleus and are 1836
times lighter.
When electrons and protons are equal in number, the atom is
neutral.
When an electron is removed, the atom will be positively
charged. An ion is an atom that has a charge on it.
Cations are positively charged atoms.
Anions are negatively atoms.

60. Isotopes
Same element, different masses and amount
of neutrons.
Nuclide = isotopic nuclei.

61. Isotopes
Carbon (atomic # 6) has three natural isotopes
with atomic weights of 12, 13 and 14.
isotope #p #n
====== == ==
C-12 6 6
C-13 6 7
C-14 6 8
Tin (Sn, atomic # 50) has ten natural isotopes with
atomic masses of 112, 114, 115, 116, 117, 118, 119,
120, 122 and 124. How many protons and neutrons
do these isotopes have?

62. Radioactive or Stable?
Radioactivity is a nuclear phenomenon: it comes as a
result of a particular structure in a nucleus.
A radioactive atom is considered unstable. All unstable
atoms emit radioactivity (usually by ejecting nuclear
particles) in order to reach a stable configuration.
This is the process of radioactive decay
So, not all atoms will be radioactive, just a small proportion
of isotopes with unstable nuclei.
The bulk of isotopes are stable, or non-radioactive.

63. Stable and Radioactive Isotopes
Carbon (atomic # 6) has three natural isotopes
with atomic weights of 12, 13 and 14.
isotope #p #n
====== == ==
C-12 6 6
C-13 6 7
C-14 6 8
C-14 is a radioactive isotope; C-12 and C-13 are stable.
Over time the proportion of C-12/C-14 and C-13/C-14
will increase until there is no C-14.
(unless some process makes new C-14...)

64. Radioactivity Inside You
Concerned about radioactivity in nature?
To keep things in perspective, consider that 0.01% of all
potassium is radioactive K-40.
Potassium is an essential element in the human body.
If your body is about 1% K, this means a 70 kg
(150 pound) person contains around
1x1021 atoms (that’s one billion trillion atoms)
of radioactive K-40.

65. Energy levels in an atom
Energy of an atom is quantized, meaning your
electrons all have discrete amounts of energy.
Electrons are thus limited to a specific energy level,
Which you learned in grade 8 was an orbital.
These main energy levels are indicated by n, and
1, 2, or 3 following it.
eg. n = 1

66. Energy levels in an atom
Each of these main energy levels are then
sub-divided into sub-energy levels, which are
indicated by the numbers s, p, d and f.
s<p<d
energy

67. Energy levels in an atom
Electrons are constantly moving, and it is impossible
to determine the position of an electron.
Experimentation has indicated the most likely area of
motion. This area is known as an orbital.
s-level = one s-orbital, two electrons
p-level = 3 p-orbitals, six electrons
d-level = 5 d-orbitals, 10 electrons

68. Shape of orbitals

69. Electron occupation of orbitals
The distribution of electrons occur according
to the Aufbau principle.
●Pauli's exclusion principle:
An orbital can carry a max. of 2 electrons, if
it spins in opposite directions.
●Electrons fill orbitals with the lowest energy
first. The atom is most stable in its lowest
state of energy.

70. Electron occupation of orbitals
●When the same kind of orbitals are
available, you first add one electron to each
orbital, and then fill them up. HUND'S RULE
●2n² = amount of electrons in that orbital

71. Aufbau diagram for 20Ca
n = 4 4s
n = 3 3p
n = 3 3s
n = 2 2p
n = 2 2s
n = 1 1s

72. ELECTRON CONFIGURATION
With electronic configuration elements are represented
numerically by the number of electrons in their shells
and number of shells. For example;
Nitrogen configuration = 2 , 5
7
N
2 in 1st shell
2 + 5 = 7
5 in 2 nd
shell
14

73. ELECTRON CONFIGURATION
Write the electronic configuration for the following
elements;
20 11 8
a) Ca b) Na c) O
23 16
40
2,8,8,2 2,8,1 2,6
17 14 5
d) Cl e) Si f) B 11
35 28
2,8,7 2,8,4 2,3

74. sp-notation
20 Ca: 1s²2s²2p63s²3p64s²
20 Ca : [Ar] 42
20 Ca 2+ : 1s22s22p63s23p6

76. Structure of the Periodic Table
GROUP
Vertical columns.
PERIODS
Seven horisontal rows.
VALENCE e- AND GROUP NUMBER
Valence e- are the same amount as the
group number.

77. Structure of the Periodic Table
PERIOD NUMBER AND ENERGY LEVELS
The period number indicates the energy level
in which the last e- are found.

78. Periodicity of the elements
Periodicity is the recurring pattern of physical
and chemical properties as you move across
the Periodic Table.

79. Periodic Law
The elements of the same group have similar
properties. These properties differ from left
to right.

80. Physical and chemical properties of the
elements are related to their atomic stucutre.

81. Atomic radius
...is the distance from the nucleus and the
outermost stable electron orbital.

82. Atomic radius left to right in a period
Atomic # increases, energy levels remain
the same.
Electrons that are added are found in the
same energy level.
Nuclear charge increases from left to right.
Greater nuclear charges causes the e- to be
really attracted to the nucleus, causing the
radius to become smaller.

83. Atomic radius top to bottom in a group
# energy levels increase from top to bottom.
Valence e- are further away from the core.
e- in the inner energy levels shield the outer
e- (screening effect)
Attractive forces of the core on the outer e-
decrease.
Atom volumes increase, so radius increase.

84. Melting and boiling points

85. Density
Forces of attraction and nature of substance
determines density.
Period – Period – non Group
metals metals
increase decrease Metals
-decrease
Non-metals -
increase
Metal bonds Giant covalent Atomic mass
Increased valence e-, structures.
stronger forces, and Regular repeating increases
tighter the atoms are pattern. Not packed
packed. closely together.

86. Ionisation energy
First ionisation energy: energy needed to
remove the 1st electron from a neutral atom in
the gaseous phase.
X(g) + ionisation energy → X+(g) + e-

87. Factors influencing ionisation energy
1. Charge of the nucleus
2. Atomic radius
3. Shielding effect
4. Repulsion forces between e- pairs

88. Factors influencing ionisation energy
1. Charge of the nucleus
The greater the atomic number, the greater
the + charge, the greater the +/- attraction.

89. Factors influencing ionisation energy
2. Atomic radius
The greater
the distance,
the smaller
the attractive
forces.
Ionisation
energy also
decrease.

90. Factors influencing ionisation energy
3. Shielding effect

91. Factors influencing ionisation energy
4. Repulsion forces between e- pairs

92. Ionisation energy - periodicity

93. Ionisation energy - periodicity
Ionisation energy from top to bottom in a
group.
Usually decreases.
Outer electrons are further away from
nucleus, shielded by inner electrons.
Increase in size, decrease in ionisation
energy.

94. Consecutive Ionisation energy
The remaining electrons are attracted to the
nucleus much stronger.
e- repulsion decreases.
Ionisation energy increases.

95. Electron affinity
Change in energy when an electron is added
to a neutral atom or ion in the gaseous
phase.

96. Electron affinity
Becomes more negative L → R.
L → R: electrons are more strongly
attracted to the nucleus.
Metals → low electron affinity. Don't take e-
easily.
Halogens → highest electron affinity. Take
e- easily.
Noble gases exempt – stable octets do not
accept e-.

97. Electron affinity
Decreases from top to bottom.
e- are further from the nucleus, weaker
force of attraction from the nucleus, electron
affinity decreases.

98. Electron negativity
The amount of energy released when an
electron is added to an atoms in the ground
state, is called electron negativity.

99. Electron negativity
Pauling scale – numbers on your periodic
table. No units
Increases from L → R in a period.
Increased radius – e- are not tightly held.
Decreases from the top → bottom in a
group. e- are held tighter by the nucleus.

100. Electron negativity

101. Electron negativity

102. Formulas of oxides
When oxygen reacts with another
element.
Metal oxides are typical crystalline ionic
solids.
High M.P. And B.P.
Dissolves in water – form bases.

103. Formulas of oxides
Ratios:
Metals –increases in each period
Non-metals – decrease in each period
Non-metal oxides – simple covalent that
dissolves in water to form acids.
Noble gases do not form oxides.

104. Formulas of halides
React with metals – form ionic solids.
React with non-metals – simple covalent
molecules.
Bonding ratio increases with metals and
decreases with non-metals.

105. Similarities in chemical properties
Read through p 121 – 124, revising the
content from your gr 8 and 9 syllabus
with your newly acquired knowledge.

106. Atomic and molecular structures

107. Atoms

108. Molecules

109. Simple molecules

110. Giant molecules - diamonds

111. Intra molecular forces

112. Intramolecular forces
Valence electrons are responsible for
bonding between different atoms.
I.F. explains the following:
How thousands of atoms stay together
in a molecule.
The microscopic properties: M.P.,
hardness, conduction &c.

113. Intramolecular forces
Covalent bonds,
ionic bonds and
metallic bonds

114. Covalent bond
Bond between non-metal and non-
metal.
Electrons are SHARED.
Forms molecules

115. Covalent bond
Hydrogen molecules

116. Covalent bond
Hydrogen chloride molecules

117. Double covalent bond
oxygen molecules

118. Double covalent bond
nitrogen molecules

119. Molecules consisting of two or more
atoms
HCl
H2O
NH3
CH4
CO2

120. Ionic bonds
Between metals an non-metals
Electrons are transferred
Form crystal lattices

121. Ionic bonds
Carries e- over
Positive and negative ions
attracted by electrostatic
forces.
This is an ionic bond
Crystal lattice that is built up of
alternative positive and negative ions.

122. Ionic bonds

123. Relative formula mass
Add the Ar of each atom.
1. Formula of ionic compound
2. Work out how many ions are present
3. Ar(element 1) + Ar (element 2)
4. No units

124. Metallic bonds
1. empty/half-filled valence orbitals
2. little energy is needed to loosen the valence
electrons – delocalised
3. atoms packed tightly – forms orderly crystal
lattice
4. valence orbital overlap
5. atoms positive charge – positive atomic
residue
6. electrostatic force between atomic residue
and sea of delocalised e-: metallic bond.

125. Metallic bonds

126. Metallic properties
PROPERTIES REASON
Metal glow Delocalised e- can reflect light –
causes shiny surface
Conduction of Delocalised e- can move freely, acts
electricity – solid/liquid as charge carriers
Conduction of heat Delocalised e- can move freely, act
as heat carriers
Malleability Atoms can slide over each other
High density Atoms packed closely together in a
metal lattice

127. Alloys
Alloys change a metals' properties. They are
usually always stronger and harder than the
pure metal.
STEEL – iron and carbon
STAINLESS STEEL – iron, chromium and
nickel
BRONZE – copper and tin