Chemistry in the Community Mid Term Review Sections 0 through 1.D - - PowerPoint PPT Presentation

Chemistry in the Community Mid Term Review Sections 0 through 1.D (p. 142 in the text book)

Section 0.C

0.3 Density of Solids and Liquids

• Density is a physical property

D = , density of water = = 1 g/cm3

• Calculate volume of a regular solid: Measure length x

width x height (units: cm3, m3, ft3 etc.)

• Measure Volume of a liquid: Use a graduated

cylinder: mL or L

• Measure volume of an irregular solid by displacement

with a graduated cylinder

• Density of water = 1.00 g/mL 1 mL = 1 cm3

mass

volume

1 g

1 cm3

¡ ¡M ¡ D ¡ ¡ ¡V ¡

Unit 1 Section A: Building Blocks of Chemistry

1.A.5 Molecular View of Matter

• Matter

– Anything that has mass and takes up space – All solids, liquids and gases are matter

• Molecules are made of atoms held together by

chemical bonds; eg. N2 , H2O

• Compounds are molecules with 2 or more different

atoms; eg. H2O

• Elements contain only one type of atom; eg. N2

– there are ~115 elements; 88 naturally occurring – most occur as atoms, some as molecules – Diatomic elements: HONClBrIF

Unit 1 A.2 Physical and Chemical Properties

• Physical properties- Properties that can

be determined without altering the chemical makeup of the material.

• Color, ductility, malleability, conductivity,

density, odor, boiling point, melting point, Freezing point, and shape

Unit 1 A.2 Physical and Chemical Properties

• Physical change in a substance does not

involve a change in the identity of the substance (grinding, cutting, melting, and boiling)

• These changes can usually be reversed

(because the identity of the substance was not changed)

Unit 1 A.2 Physical and Chemical Properties

• Chemical property relates to a

substance’s ability to undergo changes that transform it into different substances.

• Chemical change or Chemical reaction

is a change in which one or more substances are converted into different substances.

Unit 1 A.2 Physical and Chemical Properties

• A chemical change probably occurred if you
• bserve one or more of the following:
• Gas produced
• Evolution of heat and light (Burning)
• Formation of a precipitate
• Change on the surface of a solid
• Color Change to a new color that is not a result
• f the mixing of colors; Such as two clear

substances yielding a red color.

Classify as chemical or physical change

1. You make scrambled eggs. 2. You step on a piece of chalk and it becomes powdered. 3. You light a match when the electricity goes out. 4. Steam from your hot shower condenses on a cold mirror. 5. Milk turns sour. 6. Wax melts. 7. Wax burns.

1.A.5 – Particles of Matter

Modeling Matter

1. Is an element or compound shown in a? 2. Two kinds of atoms are shown in b; is this a compound? 3. Which diagrams show only molecules? Compare them. 4. Which diagrams show mixtures? Compare them.

Symbols, Formulas, and Equations 1.A.6

• Chemical Symbols represent elements

– First letter is capitalized; second letter is lower case

• Chemical Formula represent different chemical

substances, e.g. H2O

– Subscript tells how many atoms of an element

Symbols, Formulas, and Equations 1.A.6

• Chemical Formulas are the “words” in the

language of chemistry.

• Chemical equations are the sentences.
• Each chemical equation summarizes the details
• f a particular chemical reaction.
• Chemical reactions entail the breaking and

forming of chemical bonds. – Atoms are rearranged to from new

• substances. The properties of the new

substances are different from those of the

• riginal material(s)

Chemical Equation for formation of water 2 H2 + O2 à 2 H2O

Hydrogen Oxygen Water

Reactants Product

(Left side of the arrow) (Right side of the arrow)

Unit 1.A.9 The Chemical Elements Metals

• Include such elements as iron, tin, zinc, and

copper

• At room temperature they are solids
• A good conductor of heat and electricity
• Malleability- hammered or rolled into sheets
• Ductile- can be drawn in to a fine wire (high

strength)

• Luster- most metals have a grayish or silvery

luster

Unit 1.A.9 The Chemical Elements

Non Metals

• Most are gases at room temperature

(except bromine which is a liquid)

• The solid nonmetals tend to be brittle
• A poor conductor of heat or electricity

Unit 1.A.9 The Chemical Elements

Metalloids- Such as Silicon and Germanium

• Elements on the stair-step line that separate the

metals from the nonmetals

• Has some of the characteristics of metals and

nonmetals

• Are solids at room temperature
• Less malleable than metals
• Not as brittle as non-metals
• Some have a metallic luster
• Semiconductors of electricity (their ability to conduct

electricity falls between that of metals and nonmetals)

Metals, Nonmetals, Metalloids

Unit 2 A.3 Lab: Metal or Non Metal?

Mg (s) + HCl→ MgCl2(aq) + H2(g) Zn (s) + HCl (aq) → ZnCl2 (aq) + H2 (g) Fe (s) + HCl (aq) → FeCl3 (aq) + H2 (g) Sn (s) + HCl (aq) → SnCl2 (aq) + H2 (g)

How do you select a material for a specific use?

• Need suitable physical and chemical properties
• Need suitable cost

Examples:

• during WWII (1943), pennies were made of

steel, but they corroded

• Post 1982 pennies have a zinc core (cheaper)

– 97.5% Zinc

Unit 1 Section B: Periodic Trends

Periodic Table Origins 1.B.2

• By mid-1800s, about 60 elements known;

mostly solids.

• 1869 Dmitri Mendeleev published a periodic

table with elements in a regular pattern

– Elements arranged horizontally in order of increasing atomic weight (relative mass of individual atoms) – Elements with similar “combining capacity” for

• xygen and chlorine placed in vertical groups

(columns)

1.B.3 Electrical Nature of Matter

• Demonstrate charged balloon
• Like charges repel
• Unlike charges attract

+ +

• +

1.B.3 Electrical nature of Matter

• These Positive and negative charges in an

atom are:

• Protons --positive charge
• Electrons --negative charge
• Neutrons -- neutral charge
• The attraction between these particles are

the “glue” that hold the atom together

The Pattern of Atomic Numbers 1.B.4

• Properties of elements depend mostly on

the elements’ electronic structures.

• Electronic structure- The arrangement of

electrons in the atom of an element.

• Number of electrons is the same as the

number of protons for an electrically neutral atom. (not for ions)

The Pattern of Atomic Numbers 1.B.4

• The atomic number (# of protons) identifies an

element

– Every atom of carbon contains 6 protons in the nucleus

• The nucleus also contains neutrons, which have

about the same mass as protons

– The mass # of an atom = #protons+#neutrons – Electrons have 1/2000 the mass of a proton, don’t contribute to mass # – Carbon atoms can have 6, 7, or 8 neutrons

• Isotopes = atoms with same # protons, different #

neutrons

• Mass Number

Average Atomic Mass 1.B.4

Why is the atomic mass of chlorine (35.5 g) not a whole number?

Chlorine has two isotopes

76% of Cl = 35Cl (17 protons + 18 neutrons) = 35 g/mol 24% of Cl = 37Cl (17 protons + 20 neutrons) = 37 g/mol weighted average = closer to 35 g/mol = (0.76 x 35 g/mol) + (0.24 x 37 g/mol) = 35.5 g/mol

The Periodic Table

Group

1 2 3 4 5 6 7

Periods ¡ ¡ Groups ¡ ¡

Unit 1.B.7 Organization of the Periodic Table

Groups or Families

• Vertical columns of the periodic table
• Each group contains elements with similar

properties

Unit 1.B.7 The Periodic Table

Periods

• Horizontal rows
• Elements close to each other have more

similar properties than those further apart

Unit 1.B.7 The Periodic Table

Noble Gases

• Unreactive or Inert elements

Unit 1.B.7 The Periodic Table

Halogens

• Elements of group 17
• F, Cl are gases at room temperature
• Br is a reddish liquid
• I is a dark purple solid
• At is a solid
• Most reactive of the nonmetals

Unit 1.B.7 The Periodic Table

Alkali Metals

• The elements of group 1 on the periodic

table

• Silvery appearance
• Soft enough to cut with a knife
• Very reactive (usually stored in kerosene

because of their reactivity with air or moisture)

Unit 1.B.7 The Periodic Table

Alkaline Earth Metals

• Elements in group two
• Harder denser stronger than alkali metals
• Less reactive than alkali metals
• Too reactive to be found in nature as free

elements

Unit 1.B.7 The Periodic Table

Exceptions to the rule:

• Hydrogen(H) even though it’s part of

group 1, does not share the same properties of group 1

Unit 1.B.7 The Periodic Table

Transition elements

• Metals with metallic properties
• Good conductors of electricity and have

high luster

• Less reactive than alkali and alkaline earth

metals

• Do not easily form compounds
• Exist in nature as free elements

Unit 1.B.7 The Periodic Table

Inner Transition Elements

• All metals
• Lanthanides

– Shiny metals

• Actinides

– All radioactive – All beyond uranium are synthetic

Unit 1.B.7 The Periodic Table

Main Group Elements

• Groups 1-2
• Groups 13-18

Building Skills 1.B.8 Predicting Properties

• Mendeleev laid out a Periodic Table in 1869 when
• nly 60 elements were known

– Elements arranged in rows with increasing atomic weight – Elements in each column had similar chemical properties – He was so sure of his PT that he left blanks for undiscovered elements Si b.p. = 3267°C b.p. = (3267+2603)/2 = 2935°C predicted Sn b.p. = 2603°C Germanium discovered in 1886, b.p. = 2834°C

Here are formulas for several known compounds: NaI, MgCl2, CaO, Al2O3, CCl4 Knowing that elements in the same group (column) react the same ways, predict the formulas for a compound formed from: a) C and F d) Ca and Br b) Al and S e) Sr and O c) Li and Cl

1.B.10 Ions and Ionic Compounds

• Ionic compounds are made of ions (charged

particles)

– Ionic compounds are neutral; they have no net charge because the positive and negative charges

• ffset each other

– The ions are held together in crystals by attractions between the positive and negative ions

1.B.10 Formation of Ions

• Cation: an atom loses one or more electrons and

becomes positively charged

Na – 1 electron à Na+

• Anion: an atom gains one or more electrons and

becomes negatively charged

Cl + 1 electron à Cl-

+ +

1.B.11 Writing Formulas for Ionic Compounds

• 1. Write the cation first, then the anion
• 2. The correct formula contains the fewest

positive and negative ions needed to make the total electrical charge zero Give the formula for a compound containing

• a. Na+ and Cl-
• b. Mg2+ and Cl-
• c. Li+ and N3-
• d. Al3+ and O2-

Lab 1.B.12 Relative Reactivates of Metals

• Cu Mg Ag Zn
• Examples:
• Cu(s) + Mg(NO3)2 à NR
• Cu(s) + AgNO3 (aq) à Cu(NO3)2 + Ag (s)
• Order of reactants doesn’t matter.
• Order of products doesn’t matter
• Look at the 2 metals; the one in the elemental

form must be higher on the activity series than the one which is combined (in an ionic compound) for a reaction to occur.

1.B.13 Metal Reactivity

• When finely divided copper metal is heated, it

gradually reacts with oxygen in the air to produce a black copper(II) oxide:

2 Cu(s) + O2(g) à 2 CuO(s)

• Magnesium metal burns in oxygen with a blinding

light: 2 Mg(s) + O2(g) à 2 MgO(s)

• Gold does not react with oxygen
• An activity series ranks elements in order of relative

chemical reactivity. What is the relative reactivity of Cu, Mg, Au?

1.B.14 Trends in Metal Reactivity

• In general , a more reactive element

(higher on the activity series) will cause ions of a less reactive metallic element (lower in the activity series) to change to their corresponding metal.

Unit 1 Section C: Materials and Moles

Metals with highest production world- wide: 1) Iron 2) Aluminum 3) Copper U.S. used to lead production in all of

• these. This is not the

case anymore.

Unit 1.C.2 Sources and Uses of Metals

• In this section you will explore the

properties and uses of minerals and metals.

Unit 1.C.2 Sources and Uses of Metals

• Ore- a naturally occurring rock or mineral

that can be mined and from which it is profitable to extract a metal or other

• material. (Rocks with economic value)
• Rocks- Mixtures of different minerals.
• Minerals (simple definition) - compounds

which occur naturally and from which pure metals are obtained

Unit 1.C.2 Sources and Uses of Metals

• Lithosphere contains ores and minerals, which can be

mined for metals

• Many metals are not uniformly distributed on Earth
• Factors affecting whether or not to mine a metallic ore

at a specific site – Amount of useful ore at site – Percent of metal in ore – Type of mining and processing needed to extract metal – Distance of mine from refining facility and markets – Metal’s supply and demand status

Unit 2.B.1 Sources and Uses of Metals

• Copper is one of the most familiar and

widely used metal in modern society.

• Second only to silver for conductivity
• Low-cost, corrosion resistant, and ductile.
• Most common metal for electrical wiring.
• Used to produce brass, bronze, and
• ther alloys.

1.C.6 Moles: We need to count atoms

Airbags are inflated by a chemical reaction: 2 NaN3(s) 3 N2(g) + 2 Na(s)

Each airbag needs the right amount of NaN3 (sodium azide) to fill the bag with nitrogen gas

electrical

decomposition

1.C.6 Grouping items for counting

• The size of the group depends on how many items

we use at a time 1 dozen=12 gross = 12x12

1.C.6 Counting by weighing

Mass of 1 dozen objects depends on the object

1 doz jelly beans = 14 g 1 doz eggs = 680 g 1 doz grapefruit = 5400 g

1.C.6 A mole is the counting unit for atoms

1 mole = “chemists dozen”

1 pair = 2 items 1 dozen = 12 items 1 gross = 144 items 1 mole = 6.02 x 1023 items 6.02 x 1023 is “Avogadro’s number” 1 mole eggs = 6.02 x 1023 eggs 1 mole H atoms = 6.02 x 1023 H atoms 1 mole S atoms = 6.02 x 1023 S atoms

1.C.6 How big is a mole?

6.02 x 1023 grains of sand would cover New Jersey to a depth of 20 miles

1.C.6 How much does 1 mole of atoms weigh?

Since individual atoms weigh so little, chemists deal with moles of atoms. Avogadro’s number was conveniently chosen so that 1 C atom weighs 12 amu 1 mole of C atoms weighs 12 g

1.C.7 Molar mass of an element

• The periodic table gives average atomic masses of

elements; e.g. the average mass of an Fe atom is 55.8 atomic mass units (amu)

• The mass of 1 mol of Fe is 55.8 g
• The molar mass is the mass in grams of the average

atomic mass

molar mass of Fe =

55.8 g Fe mol Fe

1.C.7 Compare 1 mole samples of elements

Element # atoms in 1 mole Mass of 1 mole, grams H 6.02 x 1023 1.0 C 6.02 x 1023 12.0 O 6.02 x 1023 16.0 Fe 6.02 x 1023 55.9 Au 6.02 x 1023 197.0

1.C.9 – Composition of Minerals

l PERCENT COMPOSITION – percent by mass of

each material found in a substance.

l Ex: A penny with a mass of 2.500 grams is composed

• f 2.4375 grams of zinc and 0.0625 grams of copper.

What is the percent mass of each metal in the penny?

l % Zn = 2.4375 g Zn x 100 = 97.50% Zinc

2.500g total

• % Cu = 0.0625 g Cu x 100 = 2.50% Copper

2.500g total

l Percent composition helps geologists to describe

how much metal or mineral is present in a particular ore.

l Ex: the mineral chalcocite = source of copper metal

§ Chemical formula = Cu2S § To calculate % comp. of Copper in chalcocite:

% comp = number of moles of Cu in Cu2S x molar mass of element (Cu) molar mass of compound (Cu2S) ¡ ¡ ¡ ¡ ¡ ¡

FORMULA: ¡ ¡ ¡ ¡ ¡ ¡ ¡

% ¡comp ¡= ¡( ¡# ¡moles ¡of ¡element ¡in ¡compound ¡x ¡molar ¡mass ¡of ¡element ¡) ¡ ¡X ¡ ¡100 ¡

¡ ¡ ¡ ¡ ¡ ¡ ¡ ¡ ¡ ¡ ¡ ¡ ¡ ¡molar ¡mass ¡of ¡compound ¡ ¡

molar ¡mass ¡of ¡Cu2S ¡= ¡(2 ¡x ¡63.55 ¡g/mol) ¡+ ¡32.07g/mol ¡ ¡ ¡ ¡ ¡ ¡ ¡= ¡127.10 ¡g/mol ¡+ ¡32.07g/mol ¡ ¡ ¡ ¡ ¡ ¡ ¡= ¡159.17 ¡g/mol ¡ % ¡Cu ¡= ¡molar ¡mass ¡of ¡Cu ¡x ¡100 ¡= ¡127.10 ¡g/mol ¡x ¡100 ¡ ¡ ¡ ¡ ¡ ¡ ¡ ¡ ¡ ¡ ¡ ¡molar ¡mass ¡of ¡Cu2S ¡ ¡ ¡ ¡ ¡ ¡ ¡ ¡ ¡ ¡ ¡ ¡159.17 ¡g/mol ¡ ¡ % ¡Cu ¡= ¡0.7985 ¡x ¡100 ¡= ¡79.85 ¡% ¡Copper ¡

Mining and Refining 1.C.10 p.101

• Converting a combined metal (usually a metal

ion) in a mineral to a free metal involves a particular kind of chemical change.

• Reduction- One or more electrons are added to

an atom of an element.

• Oxidation- An electron (or more than one) is

removed from the atom of an element.

• The atom may be by itself or part of a molecule
• f a compound.

• Oxidation- An electron (or more than one)

is removed from the atom of an element.

• Example of a half-reaction representing
• xidation:
• Fe à Fe 2+ + 2e-

26 protons 26 protons 26 electrons 24 electrons (“atom”) (“ion”)

Mining and Refining 1.C.10 p.101

• Reduction- One or more electrons are

added to an atom of an element.

• Examples of a half-reactions representing

reduction:

• Cl + 1e- à Cl-
• Co2+ + 2e- à Co

Mining and Refining 1.C.10 p.101

Mining and Refining 1.C.10 p.101

• Historically “oxidation” referred to the chemical

combination of a substance with oxygen, as the term itself suggests.

• Chemists now know that in nearly all cases in

which oxygen combines with another element or compound, oxygen removes one or more electrons from the other species.

• By today's definition any reactant that causes a

species to lose one or more electrons is said to cause that species to be oxidized.

Mining and Refining 1.C.10 p.101

• Lab activity 1.B.12 p.74
• Cu(s) + AgNO3(aq) à Cu(NO3)2(aq) + Ag(s)
• Cu(s) + 2 Ag+ (aq) à Cu2+ (aq) + 2Ag(s)
• Copper Silver ion Copper (II) ion Silver metal

metal Copper atom oxidized Silver ion reduced

Mining and Refining 1.C.10 p.101

• Cu(s) + 2 Ag+ (aq) à Cu2+ (aq) + 2Ag(s)

Copper Silver ion Copper(II) ion Silver metal metal

• Also in that same lab:
• Cu 2+ (aq) + Mg(s) à Cu(s) + Mg 2+ (aq)
• Which is oxidized which is reduced?

Mining and Refining 1.C.10 p.101

• Cu 2+ (aq) + Mg(s) à Cu(s) + Mg 2+ (aq)
• Which is oxidized which is reduced?
• OIL RIG or LEO GER
• Called Oxidation-Reduction reactions or redox

reactions

• Total electrical charge on both sides of the

equation is the same.

• Charges AND atoms must balance correctly.

Mining and Refining 1.C.10

• Cu 2+ (aq) + Mg(s) à Cu(s) + Mg 2+ (aq)
• This isn’t a useful way to get copper metal

because you are “using up” another highly desirable metal…magnesium metal

• Reducing Agent- A reacting chemical

species that serves as the source of electrons.

Mining and Refining 1.C.10

Chart p. 75 (1.B.13)

• Electrometallurgy- electric current is used to

supply electrons to metal ions-reducing them.

• Pyrometallurgy- Heat is used—such as a blast

furnace---Coke(Carbon) and CO are common reducing agents

• Hydrometallurgy- Treatment of ores and other

metal containing materials by reactants in water solution.

Unit 1 Section D: Conservation and Chemical Equations

1.D Conservation

Burning coal: + à C + O2 CO2

1 carbon atom 1 oxygen 1 carbon dioxide molecule molecule

Equation with chemical formulas: C(s) + O2(g) à CO2(g) Words: One carbon atom reacts with one oxygen molecule to produce one carbon dioxide molecule

1.D.3 Example - Converting Copper

2 copper atoms + 1 oxygen à 2 copper(II) oxide molecule formula units Equation with chemical formulas: 2 Cu(s) + O2(g) à 2 CuO(s) Word equation: “2 copper atoms + 1 oxygen molecule react to produce (yields) 2 copper(II) oxide formula units”

Accounting for Atoms

C3H8 + O2 à CO2 + H2O

1 propane 1 oxygen 1 carbon dioxide 1 water molecule molecule molecule molecule

Reactant side Product side 3 C atoms 1 C atom 8 H atoms 2 H atoms 2 O atoms 3 O atoms

Polyatomic Ions

• Sodium “salts”: Na+Cl-, Na+OH- = Na+(OH)-
• Potassium salts: KCl, KOH
• Magnesium salts: MgCl2, Mg(OH)2

Mg(OH)2 = Mg2+ OH- OH- atom inventory = 1 Mg atom 2 O atoms 2 H atoms

Atom Inventory Practice

• 1. AgNO3
• 2. Sr(OH)2
• 3. Al2(CO3)3
• 4. (NH4)2S
• 5. (NH4)3PO4
• 6. Ca(HCO3)2

1.D.5 Balancing Chemical Equations

• Balancing equations is making the number
• f atoms of each element the same on the

left and right sides of the equation. (Before and after the reaction occurs)

1.D.5 Balancing Chemical Equations

• Step ¡1-­‑ ¡Determine ¡the ¡correct ¡formulas ¡of ¡the ¡

reactants ¡and ¡products ¡FIRST. ¡These ¡formulas ¡ and ¡subscripts ¡can ¡not ¡be ¡changed ¡in ¡ balancing ¡the ¡equaRon. ¡

• Example: ¡Hydrogen ¡gas ¡and ¡oxygen ¡gas ¡react ¡

to ¡produce ¡water: ¡

• H2(g) ¡+ ¡O2(g) ¡à ¡H2O ¡(g) ¡

1.D.5 Balancing Chemical Equations

• Why ¡do ¡we ¡add ¡subscripts ¡to ¡H ¡and ¡O? ¡
• H2(g) ¡+ ¡O2(g) ¡à ¡H2O ¡(g) ¡
• Diatomic ¡Elements: ¡HONClBrIF ¡
• Step ¡2-­‑ ¡Use ¡whole ¡numbers ¡(coefficients ¡in ¡

front ¡of ¡reactants ¡or ¡products) ¡to ¡balance ¡

• elements. ¡
• ¡ ¡2 ¡ ¡H2 ¡(g) ¡ ¡+ ¡__O2 ¡(g) ¡ ¡ ¡ ¡ ¡ ¡à ¡ ¡ ¡ ¡ ¡2 ¡H2O ¡(g) ¡
• Balances ¡ ¡ ¡ ¡ ¡ ¡Stays ¡1 ¡ ¡ ¡ ¡ ¡ ¡ ¡ ¡ ¡ ¡Balances ¡Oxygen ¡ ¡ ¡ ¡ ¡

1.D.5 Balancing Chemical Equations

• Balance ¡1 ¡element ¡at ¡a ¡Rme, ¡going ¡back ¡and ¡

forth ¡between ¡sides ¡beginning ¡with ¡the ¡ element ¡furthest ¡to ¡the ¡lea. ¡

• Step ¡3-­‑ ¡Check ¡the ¡number ¡of ¡atoms ¡of ¡

elements ¡on ¡both ¡sides. ¡ ¡

1.D.5 Balancing Chemical Equations

Helpful information:

• For elements occurring more than once on

either the left or the right --leave those elements to last.

• Balance uncombined elements last.
• If a polyatomic ion (example: NO3
• ) is on

both sides, balance the whole ion, not each element in it.

2 Al (s) + 3 Pb(NO3)2 (aq) à 2 Al(NO3)3 (aq) + 3 Pb (s)

1.D.5 Balancing Chemical Equations

• If there are 2 of an atom (or ion) on one

side and 3 of the same atom or ion on the

• ther side, make them both 6 by using a

coefficient of 3 in front of the first and 2 in front of the second.

1.D.5 Balancing Chemical Equations

• If ¡H2O ¡is ¡on ¡one ¡side ¡and ¡OH-­‑ ¡is ¡on ¡the ¡other ¡

side ¡(by ¡itself ¡or ¡as ¡part ¡of ¡an ¡ionic ¡ compound), ¡Rewrite ¡H2O ¡as ¡HOH ¡

• Example: ¡ ¡

Na(s) ¡+ ¡H2O(l)à ¡NaOH(aq) ¡+ ¡H2(g) ¡

• 2Na(s) ¡+ ¡2 ¡HOH(l) ¡ ¡à ¡2 ¡NaOH(aq) ¡+ ¡H2(g) ¡

¡

1.D.5 Balancing Chemical Equations

• A balanced chemical equation represents

the law of__________ ___ _____

• Conservation of Matter

Calculations with molar mass

moles to grams

• A Si chip contains 0.05 mole Si. How many g Si are

in the chip?

• molar mass of Si =

Set up the question: 0.05 mole Si = ? g Si

0.05 mole Si x 28.1 g Si = 1.41 g Si mole Si

28.1 g Si mole Si

Calculations with molar mass

grams to moles

• An iron bar is 16.8 g. How many moles Fe are in the

sample?

• molar mass of Fe =

Set up the question: 16.8 g Fe = ? atoms Fe

16.8 g Fe x 1 mole Fe = 0.30 moles Fe 55.8 g Fe

55.8 g Fe mole Fe

l If ¡you ¡need ¡to ¡make ¡25 ¡grams ¡of ¡H2O, ¡how ¡many ¡

grams ¡of ¡PbO ¡do ¡you ¡need ¡to ¡start ¡with? ¡

l This ¡can ¡be ¡determined ¡by ¡using ¡a ¡conversion ¡

factor… ¡

l Conversion ¡Factor ¡– ¡Relates ¡the ¡amount ¡of ¡one ¡

substance ¡in ¡a ¡chemical ¡reacRon ¡to ¡another ¡ substance ¡in ¡the ¡equaRon ¡

l Allows ¡us ¡to ¡convert ¡from ¡one ¡unit ¡to ¡another, ¡or ¡one ¡

substance ¡in ¡the ¡reacRon ¡into ¡another. ¡

Conversion with a familiar concept...

How many minutes are in a year?

What information do you need to solve this? Write out the conversions that you will do.

Grams ¡A ¡à ¡Moles ¡A ¡à ¡Moles ¡B ¡à ¡Grams ¡B ¡

Mole ¡Map ¡

A ¡“map” ¡for ¡sejng ¡up ¡conversion ¡factors ¡from ¡ ¡ chemical ¡reacRons ¡

1.D.9 Conservation in the Community

• The Earth is like a spaceship: the resources

“on board” are all that are available

– Renewable resources can be replenished by natural processes

• water, air, fertile soil, plants, animals
• as long as natural cycles are not disturbed

– Nonrenewable resources cannot be readily replenished

• metals, natural gas, coal, petroleum
• we are using petroleum faster than nature is making it

1.D.9 How can we avoid depleting nonrenewable resources?

• 1. Use less (“source reduction”)
• 2. Find substitute materials with similar

properties; refurbish or reuse (e.g. used car parts, printer cartridges)

• 3. Recycle