Physical Science I

1
Unit II

• Physical Science I
• Chemistry

2
Chemistry Terms

• Matter - anything that has mass and takes up
  space
• Chemistry - is the study of matter, its
  properties and the changes or chemical reactions
  that matter can undergo.
• Ex. rusting, combustion of fuel/candle wax,
  explosion of TNT, vinegar and baking soda

3

• Pure Chemistry - describing known substances and
  discovering new compounds for research purposes.
• Applied Chemistry the search for uses for
  chemical substances. Modern society demands
  chemistry understanding.
• Ex. technology government.

4

• Mass- The amount of matter an object contains in
  grams (g).
• States of Matter 3 physical states
• Solid definite volume/shape
• Liquid definite volume/indefinite shape
• Gas indefinite volume/shape
• (aq)aqueous dissolved in H2O

5

• Physical Property- a characteristic of a
  substance can be observed without changing into
  a new substance.
• Ex. State of matter, hardness, melting point,
  boiling point, odour, solubility colour,
  malleability, ductility, brittleness,
  conductivity.

6

• Physical Change- a change in state of a substance
  (no new substance formed).
• Ex. melting (s to l) evaporation/boiling (l to
  g) condensation (g to l) sublimation (s
  to g)
• H2O (s) H2O (l)

heat
7

• Chemical Property- a characteristic behaviour of
  a substance that occurs when a substance changes
  into a new substance.
• Ex.
• 2 Mg(s) O2(g) 2 MgO(s) Mg ribbon

light energy
8

• Chemical Change- a change in which one or more
  NEW substances are formed.
• Ex.
• coal combustion
• C(s) O2(g) 2 CO2(g) heat
• rusting
• 4 Fe(s) 3 O2(g) Fe2O3(g)

9

• Indicators of Chemical Change
• new colour
• heat/light given off
• Bubbles of gas
• Precipitate (solid) formation
• Change is difficult to reverse

10

• Mixture contains 2 or more pure substances.
• Two Types
• Homogeneous Mixture- aka solution- have only
  one visible phase throughout.
• ex. air, apple juice, salt water
• Heterogeneous Mixture- contain 2 or more visible
  components or phases
• ex. soil, soup

11

• Pure Substance made up of only one type of atom
  or atom combination.
• (Ex. O2, H2O)
• Stays the same in response to physical change.
• Two Types Compounds and Elements

12

• Compounds - pure substances that contain two or
  more different elements in a fixed proportion
• ie., CO2, H2O, C6H12O6 and NaCl
• Can be broken down to elements via chemical
  means.
• Ex. 2 NaCl (l) 2 Na (l) Cl2 (g)

electricity
13

• Elements - pure substances that CANNOT be broken
  down into simpler substances by regular
  laboratory conditions made up of 1 type of atom.
• ie., oxygen, nitrogen, carbon and phosphorus
• Element Symbols are always written with the first
  letter uppercase and the second letter lowercase.
• Ex. Au, Mg, Ar
• Element Names are always written in lowercase
  letters.

14

• Diatomic Molecules There are 7 elements that
  are diatomic gases in their natural state.
• These are H2 O2 F2 Br2 I2 N2 Cl2
• Also P4 and S8
• How can we remember these?
• HOFBrINCl PS! (or an upside down L on the
  periodic table).

15
Matter Flow Chart
(Ionic, Molecular, Acids)
(Metals, Nonmetals)
(solutions)
16

• Reactants starting materials.
• Products new substances formed.
• Chemical Reaction
• Reactants Products

go to form
17
Periodic Table

• Mid 1800s- 65 known elements.
• Began to recognize patterns after recording
  reactivity, masses, etc.

18

• Dmitri Mendeleev (1834-1907)
• Wrote out elements in order of increasing atomic
  mass, result was a table.
• periodic table- periodic meaning repeating
  patterns and properties.
• We now organize the periodic table according to
  atomic number
• Also organized according to of electrons (e-)
  in atoms of each element.

19
Periodic Table- A Review

• It is designed to arrange elements in a pattern
  that helps us predict properties and bonding
  patterns of elements.
• Elements are organized by Atomic Number and
  Number of Electrons.
• The periodic Table is arranged in rows and
  columns.

20

• Period horizontal row (7 in total)
• atomic mass and atomic number increase( )from
  L to R.
• Group/Family vertical columns (18)
• Elements of the same group have similar
  but not identical properties.
• Some groups have species names
• Group 1- Alkali Metals Group 17- Halogens
• Group 2- Alkaline Earths Group 18- Noble
  Gases
• Lanthanides (rare earth)
• Actinides

21

• Groups have 2 numbering systems
• New Group 1-18
• Old Roman Numerals/Letters
• IA-VIIIA - Representative Elements
• IB-VIIIB - Transition Elements.

22
Metals and Nonmetals
He 2
H 1
1
C 6
Li 3
N 7
O 8
F 9
Ne 10
B 5
Be 4
Nonmetals
2
Na 11
Al 13
Si 14
P 15
S 16
Cl 17
Ar 18
Mg 12
3
K 19
Ca 20
Sc 21
Ti 22
V 23
Cr 24
Mn 25
Fe 26
Co 27
Ni 28
Cu 29
Zn 30
Ga 31
Ge 32
As 33
Se 34
Br 35
Kr 36
4
METALS
Rb 37
Sr 38
Y 39
Zr 40
Nb 41
Mo 42
Tc 43
Ru 44
Rh 45
Pd 46
Ag 47
Cd 48
In 49
Sn 50
Sb 51
Te 52
I 53
Xe 54
5
Cs 55
Ba 56
Hf 72
Ta 73
W 74
Re 75
Os 76
Ir 77
Pt 78
Au 79
Hg 80
Tl 81
Pb 82
Bi 83
Po 84
At 85
Rn 86
6

Fr 87
Ra 88
Rf 104
Db 105
Sg 106
Bh 107
Hs 108
Mt 109
7
W
Ce 58
Pr 59
Nd 60
Pm 61
Sm 62
Eu 63
Gd 64
Tb 65
Dy 66
Ho 67
Er 68
Tm 69
Yb 70
Lu 71
La 57
Th 90
Pa 91
U 92
Np 93
Pu 94
Am 95
Cm 96
Bk 97
Cf 98
Es 99
Fm 100
Md 101
No 102
Lr 103
Ac 89
23
Types of Elements The staircase line divides the
elements into two major categories the metals
and the nonmetals The ratio of metals to
nonmetals is about 41 The Metals Metals are
shiny, electrically conductive elements. They
are also malleable (can be hammered into shapes)
and ductile (can be stretched into wire). With
the exception of mercury, they are all solids at
room temperature (25C).
24
The Nonmetals

• Nonmetals are dull and are very poor conductors
  or nonconductors of electricity.
• The solid nonmetals are brittle.
• As a group, the nonmetals exhibit the three
  states of matter at room temperature.
• eg, carbon is a solid, nitrogen is a gas, and
  bromine is a liquid.

25

• The location of hydrogen in the periodic table is
  unusual.
• Hydrogen is a nonmetal, but in some periodic
  tables it is located in the top left hand corner
  of the periodic table (i.e. on the metals side).
• Due to the fact that hydrogen has some metallic
  properties in addition to nonmetallic properties.
• 2 groups Alkali Metals and Halogens (1 and 17).

26
The Metalloids

• The metalloids are elements that possess both
  metallic and nonmetallic properties.
• For example, silicon is shiny and conducts
  electricity (like a metal) , but it is brittle
  (like a nonmetal).
• Metalloids are also known as the semimetals.

27
Representative Elements

• The representative elements illustrate the
  entire range of the properties of the elements.
• Group 1, 2, 13-18.
• Sometimes known as the group A elements, they are
  organized into chemical families based on their
  specific chemical and physical properties.

28
The Transition Elements

• The transition elements are all metals.
• They are different from the representative
  elements because of their electron arrangements
  which in turn gives them properties that are a
  little different from the metallic representative
  elements.
• Inner Transition Elements
• (Lanthanides, Actinides) Same as transition, but
  removed from main table as a matter of
  convenience in organizing table.

29

• Elements in the periodic table are organized
  based on shared properties.

Metalloids (staircase)
30
Molecular Substances

• Exist as groups of atoms called molecules
• Molecules are substances composed of nonmetallic
  elements

Nitrogen N2
Methane CH4
31
You must memorize these

• Mono-atomic molecular elements

Noble Gases (group VIIIA or 18)
He Helium
Ne Neon
Ar Argon
Kr Krypton
Xe Xenon
Rn Radon
32
Do you know this one?
neon
Ne
33
Diatomic molecular Elements

• All have two identical atoms

eg

O
I2
iodine
34
How about this one?
nitrogen
N2
35
Memorize the diatomic molecular elements
Hydrogen H2
Oxygen O2
Fluorine F2
Bromine Br2
Iodine l2
Nitrogen N2
Chlorine Cl2
Just remember the famous chemist Dr. HOFBrINCl
36
remember the
Molecular Elements
ozone O3
Sulfur S8
Phosphorus (red) P4
Phosphorus (white) P10
37
Molecular Compounds

• Consist of two or more nonmetallic elements
• 2 types of molecular compounds
• Binary molecular compounds
• Ternary molecular compounds

38
Binary vs ternary Molecular Compounds
Type of Compound Molecular formula name
Binary H2O water
Binary H2O2 hydrogen peroxide
Binary NH3 ammonia
Binary CH4 methane
Ternary CH3OH methanol
Ternary C2H5OH ethanol
Ternary C12H22O11 sucrose
Memorize the names and formulas of common
molecular substances as per the chemistry facts
sheet
39
Binary Molecular Compounds

• Composed of 2 nonmetals
• CO2 , CCl4 , BF3 are examples
• Many are identified by common names
• ie., water H2O ammonia NH3
• System for naming and writing formulas
  established by I.U.P.A.C.
• International Union for Pure and Applied
  Chemistry

40

• Requires a system of prefixes

Number prefix
1 mono
2 di
3 Tri
4 Tetra
5 Penta
6 hexa
7 Hepta
8 Octa
9 nona
10 deca
This table is on your chemistry facts page
41
RULES FOR NAMING BINARY MOLECULAR COMPOUNDS

• Write the name of the first element of the
  formula in full.
• Shorten the name of the second element and add
  the ide ending.
• Use prefixes to indicate the number of atoms of
  each element in the molecular formula.
• The prefix mono on the first name is optional.

42

• Write the IUPAC name for CCl4
• The first element is C.
• Its full name is carbon.
• The second element is chlorine.
• Its name is shortened to chlor, and the suffix
  ide is added to give chloride.
• The prefix mono (1) is added to carbon, and the
  prefix tetra (4) is added to chloride to give the
  name
• monocarbon tetrachloride.
• The prefix mono can be omitted from the first
  element name to give
• carbon tetrachloride.

43

• Name N2O4
• Nitrogen oxide
• Dinitrogen tetraoxide (tetroxide)

44
U do

• Name
• B2H6
• Diboron hexahydride

45
Chemical Bonding

• Molecular compounds like B2H6 are held together
  by bonds.
• A chemical bond is the force of attraction
  between atoms.
• In molecular compounds the bond is the force of
  attraction occurs between nonmetallic elements
• This type of bond is called a covalent bond

46
Atomic Theory

• Atom the smallest particle of an element that
  retains the properties of that element.
• Atoms are thought to be composed of negatively
  charged particles called electrons and a dense
  central region called the nucleus

47

• Within the nucleus are found positively charged
  particles called protons and neutral particles
  known as neutrons
• Electrons are believed to exist a specific
  distances from the nucleus called energy levels

48
Electron Energy Level Diagrams Representative
Elements

• Show electron arrangements within an atoms
  energy levels.
• We can predict them using the following
• Atomic number
• Period number
• Group number
• Electrons per energy level

49
Electron Energy Level Diagrams Representative
Elements

• Atomic number
• Atomic represents the number of protons in the
  nucleus of an atom.
• the protons electrons. 
• Example Carbon has atomic 6 which means C has
  6 protons and 6 electrons)

nucleus
6
6 electrons
50
Electron Energy Level Diagrams Representative
Elements

• Period Number
• tells us how many energy levels contain electrons
• Eg Carbon is in the second row of the periodic
  table thus it is in period 2 and has electrons in
  2 energy levels.

2nd
1st
6
51
Electron Energy Level Diagrams

• Group (family number)
• The group tells us about electrons in the outer
  energy level of an atom (valence electrons)

52
Electron Energy Level Diagrams Representative
Elements

• We will deal with the representative elements
  only
• (for groups 13 and above valence electrons the
  last digit)
• 13, 14, 15, 16, 17, 18
• C has 4 valence electrons since it is group 14

2nd
4e-
1st
6
53
Electron Energy Level Diagrams Representative
Elements
Period of elements in the period Energy level max of e-
1 2 1 2
2 8 2 8
3 8 3 8
4 18 4 18
54
Electron Energy Level Diagrams Representative
Elements
2nd
4e-
2e-
1st
6
55
Electron Energy Level Diagrams Representative
Elements

• Example Draw an electron energy level diagram
  for an atom of aluminum

Group 13
3e-
Fill in the nonvalence electrons with the
maximums per energy level
Period 3
8e-
2e-
Atomic number 13
13
56
Energy level diagrams - ions

• Recall the noble gases are unreactive
• All noble gases have filled valence energy levels

57

• atoms react by changing the number of electrons
  to try and get the same structure of the nearest
  noble gas
• In other words, atoms either gain or lose
  electrons to become stable
• Metals lose electrons to have the same electron
  arrangement as the nearest noble gas
• Nonmetals gain electrons to have an electron
  arrangement of the nearest noble gas.

58
(No Transcript)
59
(No Transcript)
60
Ions

• Metals lose electrons to form positive ions
  called cations
• Nonmetals gain electrons to form negative ions
  called anions
• Both ions have noble gas stability

61
Ionic compounds

• Composed of oppositely charged ions
• Held together by ionic bonds
• 3 categories of ionic compounds
• 1) Binary ionic compounds
• simple ions (only single charges)
• multivalent ions (more than one charge)
• 2) Polyatomic ions (complex ions)
• 3) Hydrates

62
Binary ionic compounds

• Binary ionic compounds are composed of a metal
  ion ()  and non-metal ion (-).
• Naming binary ionic compounds
• Name the cation () by writing the full name of
  the metal.
• Name the anion (-) by shortening the name of the
  element and add the -ide ending.

63
Binary ionic compounds

• NaCl     sodium and chlorine  
• sodium chloride
• CaF2  calcium and fluorine     
• calcium fluoride
• K2O  potassium and oxygen  
• potassium oxide
• IMPORTANT Do Not use prefixes - they are for
  molecular compounds (two non-metals)

64
Rules for Writing Binary Ionic Formulas

• Write down the symbols of the ions involved.
• Determine the lowest whole number ratio of ions
  that will give a net charge of zero.
• Write the formula removing all charges.

65
Write a chemical formula for a compound that
contains Calcium ions and Bromide ions.
Sample

• Write down the symbols of the ions involved
• Calcium is group IIA, Ca2
• Bromide is group VIIA, Br
• Determine the lowest whole number ratio of ions
  that will give a net charge of zero. Use the
  crossover method
• Ca2 Br
• Ca Br2

66
Binary Ionic CompoundsThe Stock System

• Ions of a certain elements can have more than one
  possible charge.
• Such elements are called multivalent species.
• Example 1 tin forms two common ions
• Sn2 and Sn4
• The Stock System is used to name ions like these
• Sn2 is called tin (II) and Sn4 is called tin
  (IV)

67
Stock System

• Example too! Cu is copper (I)
• Cu2 is copper (II)
• The periodic table lists the ions that have stock
  names.

68
Ion Stock name
Cu Copper (I)
Cu2 Copper (II)
Fe2 Iron (II)
Fe3 Iron (III)
Sn2 Tin (II)
Sn4 Tin (IV)
Pb2 Lead (II)
Pb4 Lead (IV)
69
Stock System

• Egg Sample Write the chemical formula for
  iron(II) chloride.
• Write the symbols of the ions involved
• Iron (II) (the roman numeral tells us it has a 2
  charge) Fe 2
• Chloride (has a 1- charge)  Cl
• Determine the lowest whole number ratio of ions
  that will give a net charge of zero.
• Use the criss cross method
• Fe 2 Cl
• FeCl2

70

• Write formulas for
• Titanium (IV) fluoride
• Titanium (II) fluoride
• Nickel (II) oxide
• Lead (IV) sulfide

71
Naming Ionic Compounds - Polyatomic ions

• polyatomic ion (complex ion) - is a group of
  atoms that are covalently bonded which then gain
  or lose electrons to become stable
• Example
• The ammonium ion, NH4, consists of one nitrogen
  atom and four hydrogen atoms which as a group
  have lost one electron.

72
Table of Complex Ions
73
Writing Chemical Formulas for Compounds with
Polyatomic Ions

• write the cation symbol first and the anion
  symbol last. 
• balance the charges by providing the appropriate
  numerical subscript for each ion. 

74

• Write the a chemical formula for each compound
• magnesium chlorate
• Mg2 ClO3-
• put brackets around the complex ion
• Mg2 (ClO3)
• criss cross the charges
• Mg (ClO3)2
• iron(III) sulfate
• Fe2(SO4)3

75
Naming Ionic Compounds Ionic Hydrates

• An ionic hydrate is a compound that decomposes
  upon heating to release water
• Water is part of its crystalline structure.

CuSO4?5H2O is copper(II)sulfate pentahydrate
76
Each ionic hydrate has two parts to its name
A number of molecules of water
Ionic salt
CoCl2
2 H2O
?
A separator

• cobalt (II) chloride

dihydrate
77
U do

• zinc sulfate heptahydrate
• potassium sulfate decahydrate
• nickel (II) nitrate tetrahydrate

78
Acids hydrogen compounds

• All are hydrogen compounds dissolved in water
• Acids can be simply defined as substances that
  release hydrogen ions (H) in water
• Substances dissolved in water are denoted by a
  subscript (aq) written after their formula
• Eg. HCl(aq)
• Acids turn blue litmus red

79
Rules for Naming Acids

• name the hydrogen compound as if it were an ionic
  compound.
• (all of these compounds should end in - ide,
  -ate, or -ite.)
• depending on the ending convert the ionic name to
  the acid name.
• Ionic name acid name
• hydrogen _______ide  ?  hydro _____ic acid
• hydrogen _______ate  ?   _________ ic acid
• hydrogen _______ite   ?   ________ous acid

80
Rule 1 hydrogenide

• If the aqueous hydrogen compound begins with
  hydrogen and ends in ide, then
• Replace hydrogen with hydro and ..
• replace the ide ending of the anion with ic
  acid
• Example HCl (aq) 
• Hydrogen chloride ?  hydrochloric acid

81
Rule 2 hydrogen..ate

• If the aqueous hydrogen compound begins in
  hydrogen and ends in ate then
• drop the name hydrogen (do not replace it)
• replace the ate ending of the anion with -ic
  acid
• Example HClO 3 (aq) 
• hydrogen chlorate   ?  chloric acid

82
Rule 3 hydrogenite

• If the aqueous hydrogen compound begins in
  hydrogen and ends in ite then
• drop the name hydrogen (do not replace it)
• replace the ite ending with -ous acid
• Example HNO2 (aq) 
• hydrogen nitrite   ?  nitrous acid

83
Bases

• Bases are substances that behave in opposition to
  acids.
• ionic compounds that contain the hydroxide ion
  (OH-).
• Sodium hydroxide NaOH (aq)

Sample
84
Properties of Bases

• turn red litmus blue
• neutralize acids
• have high pH (gt 7)
• form slippery solutions
• tend to have a bitter taste

pH scale
85
Chemical Change

• communicated in sentence form or as chemical
  equations.

86
Chemical equations have four parts

• 1 chemical formulas
• 2 subscripts for states of matter
• (s) solid
• (l) liquid
• (g) gas
• (aq) aqueous - dissolved in water
• 3 numerical coefficients
• indicates how many atoms/molecules are involved
• 4 reaction symbols
• the "" sign on the reactants (left) side is read
  as "reacts with"
• the arrow ( ? ) is read as "to produce"
• the "" sign on the products (right) side is read
  as "along with".

87
Chemical Equations

• Two molecules of diesel fuel react with 49
  molecules of oxygen to produce 32 molecules of
  carbon dioxide and 34 molecules of water.

88
Evidence for Chemical Change

• Chemical changes involve changes in make up - new
  substances are formed with new properties
• Physical changes involve changes in state without
  a change in make up.

89
Evidence of Chemical Change

• 4 indicators of chemical reaction
• energy change
• colour change
• precipitate formation
• gas formation

90

1. a rock warmed by the sun all day loses its heat
   at night
2. milk goes sour when left out of the fridge
3. bubbles form in a glass of cold water as it warms
   
4. bubbles and steam rise out of a kettle of boiling
   water
5. paint dries on a hot day.

91
The Law of Conservation of Mass

• In a chemical reaction the mass of the reactants
  before a chemical reaction equals the mass of the
  products after the reaction is complete.
• Antoine Lavoisier
• placed mercury(II) oxide powder (a red powder) in
  a test tube, sealed it, and then weighed it
  carefully

92

• heated it and observed that the red powder
  gradually changed into a grey liquid
• reweighed the sealed tube after the reaction was
  complete and observed that its mass had not
  changed
• opened the tube and noticed a rapid release of a
  gas which was later learned to be oxygen. The
  grey liquid was mercury metal.

93
Balancing Chemical Equations

• All representations of a chemical change must
  reflect the law of conservation of mass
• Chemical equations must obey the law of
  conservation of mass
• The number of atoms of each element must be the
  same on both sides of the equation
• For this reason all chemical equations we write
  must be balanced.

94
Balancing Chemical Equations

• Mercury(II)oxide forms mercury and oxygen
• Reactants Products
• HgO(s) ? Hg(l) O2(g)
• 1 Hg 1 Hg
• 1O 2 O equation is
  unbalanced
• 2HgO(s) ? 2 Hg(l) O2(g)
• 2 Hg 2 Hg
• 2O 2 O equation is
  balanced
• 2HgO(s) ? 2 Hg(l) O2(g)