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Physical Science I Chemistry

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1 Physical Science I Chemistry
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) 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) CO2(g) + heat rusting 4 Fe(s) 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 Can be broken down to elements via chemical means.
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) Na (l) + Cl2 (g) electricity

13 Element Names are always written in lowercase letters.
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 How can we remember these?
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)
Mixture Heterogeneous Homogeneous Pure Substance Compound Element atom (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 ( ) 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 Metals and Nonmetals Nonmetals Metalloids H He Li Be B C N O F
1 He 2 1 Li 3 Be 4 B 5 C 6 Nonmetals N 7 O 8 F 9 Ne 10 2 Na 11 Mg 12 Al 13 Si 14 P 15 S 16 Cl 17 Ar 18 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 Metalloids 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 La 57 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 Ac 89 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

23 Types of Elements The Metals
The staircase line divides the elements into two major categories: the metals and the nonmetals The ratio of metals to nonmetals is about 4:1 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 (25°C).

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 You must memorize these

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)
(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 H2O2 hydrogen peroxide NH3 ammonia CH4 methane Ternary CH3OH methanol C2H5OH ethanol 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: prefix 1 mono 2 di 3 Tri 4 Tetra 5
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

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 monocarbon tetrachloride.
Sample 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 Sample 2 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 atom’s 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 4e- 2nd 1st 6+

53 Electron Energy Level Diagrams Representative Elements
Period #of elements in the period Energy level max # of e- 1 2 8 3 4 18

54 Electron Energy Level Diagrams Representative Elements
2nd 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.



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 Ca Br – Ca Br2

66 Binary Ionic Compounds The 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 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. sample 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: hydrogen…ide 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: hydrogen…ite 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 Sample 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 (> 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 a rock warmed by the sun all day loses its heat at night
milk goes sour when left out of the fridge bubbles form in a glass of cold water as it warms bubbles and steam rise out of a kettle of boiling water 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
sample Mercury(II)oxide forms mercury and oxygen Reactants Products HgO(s) → Hg(l) O2(g) 1 Hg Hg 1O O equation is unbalanced 2HgO(s) → 2 Hg(l) O2(g) 2 Hg Hg 2O O equation is balanced

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