Energy and Matter in Chemical Change

Science 10 PreAP Energy and Matter in Chemical Change

Understanding Matter A. WHMIS  workplace hazardous materials information system  all chemicals are treated with respect  WHMIS has been developed to provide guidelines for handling, storage and disposal of reactive materials

compressed gas corrosive flammable and combustible

poisonous and poisonous and infectious material oxidizing material infectious causing causing immediate and other toxic effects serious toxic effects

dangerously biohazardous reactive material infectious material

B. Properties  physical properties are properties you can see and measure eg) colour, state, density, ductility, malleability, boiling point  chemical properties are properties used to describe how substances will react with each other eg) combustion, rusting, decomposition C. Classification of Matter

All Matter

Mixtures Pure Substances

Homogeneous Heterogeneous Elements Compounds

Alloys Colloids Metals Ionic

Suspension Non-Metals Molecular Solutions

Mechanical Metalloids Mixture

 homogeneous mixture is a mixture of 2 or more substances that has uniform properties (that appear as one) eg) salt water, kool-aid, air

 an alloy is a homogeneous mixture of 2 or more metals eg) brass = copper + zinc steel = iron + chromium + carbon

 heterogeneous mixture is a mixture of 2 or more substances and individual components are visible eg) cereal, chocolate chip cookies, bread

 suspension is a mechanical mixture in which the components are in different states eg) mud

 colloids are mechanical mixtures in which the suspended substances cannot be easily separated from the other substances eg) milk

 elements are substances (metals, non-metals or metalloids) that are pure and cannot be broken apart eg) hydrogen, helium, carbon etc.  compounds are two or more elements combined; cannot be separated by physical means eg) water, carbon dioxide etc.  we will be focusing on elements and compounds in this course!

D. Atomic Structure  atoms are the building blocks of ALL matter  consist of a tiny nucleus and a huge “cloud” region eg) ball bearing in middle of football field

nucleus

cloud region

 nucleus makes up 99% of the mass of an atom  cloud region makes up most of the volume of an atom

Subatomic Particles 1. Proton (p + )  positive charge  found in nucleus  determines the type of element

2. Neutron (n 0)  no charge  found in nucleus  used to hold nucleus together

3. Electron ( e  )  smallest subatomic particle  negative charge  found in “cloud” region  arranged in energy levels  maximum # of electrons in each level (for first 20 elements only!): Level 1 = 2 e Level 2 = 8 e Level 3 = 8 e Level 4 = we will only go up to 2 e Atomic Mass  atomic mass unit (amu) is the mass of a proton or neutron = 1.7  1024 g  net charge for all atoms is zero  # e = # p+  mass number is the sum of the protons and neutrons (rounded to the nearest whole number)  used to find the number of neutrons eg) lithium atomic number = 3 atomic mass = 6.94 = 7 # protons = 3 # neutrons = 7 – 3 = 4

Isotopes  atoms that have the same number of protons but a different number of neutrons  atomic mass on the periodic table is an average mass based on the percentage abundances of all naturally occurring isotopes of the element

isotope notation: A X = symbol X A = mass # (#p+ + n°) Z Z = atomic # (#p+)

Examples copper - 64

64 # p+ = 29 Cu # e- = 29 29 #n° = 64 - 29 = 35

copper - 62

62 # p+ = 29 Cu # e- = 29 29 #n° = 62 - 29 = 33 Your Assignment: pg 1

E. Atomic Theory 1. Dalton’s Atomic Theory: 1808  all matter is composed of tiny, indivisible particles called atoms  atoms of an element have identical properties  atoms of different elements have different properties  atoms of two or more elements can combine in constant ratio to form new substances eg) H:O ratio 2:1 → H2O → water H:O ratio 2:2 → H2O2 → hydrogen peroxide

2. J.J. Thompson: 1897  credited with discovery of electrons  “raisin bun” model or “plum pudding” model  atom is a sphere which is positive, with negative electrons embedded in it like raisins in a bun  most of the mass is associated with the positive charge

3. Ernest Rutherford: 1911  atoms have a nucleus which is positive and has most of the mass  most of the atom is empty space occupied by the moving negatively charged electrons  proposed the existence of protons

4. Neils Bohr: 1913  electrons move in circular orbits around the nucleus  cannot exist between orbits

5. James Chadwick: 1932  showed that the nucleus must contain heavy neutral particles to account for all of the atom’s mass (neutrons)

6. Schrodinger/de Broglie: 1930  quantum mechanical model  electrons have distinct energy levels  exact locations of electrons are not defined, but the probable location in a region of space can be predicted F. The Periodic Table  the periodic table was developed by Dmitri Mendeleev in the mid 1800’s http://www.chemicool.com/

1. Atomic Number  number of protons in one atom of an element  increases from left to right and top to bottom

2. Properties  3 major categories: 1. metals are good conductors, strong, malleable (pound into thin sheet), ductile (can draw into a wire, bendable), have high luster; are found on left side of stair case 2. non metals are poor conductors, non-lustrous, weak, etc… opposite properties to metals; found on right side of staircase 3. metalloids show properties of both metals and nonmetals; found above and below staircase

3. Groups  18 vertical columns are called groups or families  2 labelling systems: Roman Numerals with letters or ordinary numbers IA, IIA, IIIB, IV, V, VI, VII, VIII etc 1, 2, 3, 4, 5, 6, 7, 8

 Group 1 (IA) Alkali Metals  Group 2 (IIA) Alkaline Earth Metals  Group 17 (VIIA) Halogens  Group 18 (VIIIA) Noble (Inert) Gases  Lanthanide Series (57-71) Rare Earths  Groups 3-12 (B series) Transition Elements  elements in each group share similar chemical properties (reactivity) although intensity changes  reactivity increases down group for metals and up group for nonmetals  group number indicates how many electrons are in the outermost energy level

4. Periods  horizontal rows  show a trend in reactivity which changes from left to right  each time you move to a new period you start the trend over

G. Electron Energy Level Representations (EELR)  nucleus – shows # p+ and n0  energy levels – shows # of e in each level (# of levels = period element is in)  valence electrons are the e in outermost energy level (same as group #, ignoring the 1 in front of groups 13-18 )

Examples

sodium - atomic # 11 -11 mass # 22.99 = 23 # p+ = 11 # e- = 23 - 11 = 12 #n° = 1 e- 8 e- 11 e- 2 e- p+ = 11 n° = 12 argon - atomic # 18 - mass # 39.948 = 40 # p+ = 18 # e- = 18 #n° = 40 – 18 = 22

8 e- 18 e- 8 e- 2 e- p+ = 18 n° = 22 Draw the EELR for the following: 1. potassium 2. chlorine 3. beryllium 4. calcium potassium - atomic # 19 chlorine - atomic # 17 - mass # 39.10 = 39 - mass # 35.453 = 35 # p+ = 19 17# p+ = # e- = 19 17# e- = #n° = 39 – 19 = 20 35#n° – 17 = = 18

1 e- 8 e- 7 e- 19 e- 8 e- 8 e- 17 e-

2 e- + 2 e- p+ = 19 p = 17 n° = n° = 20 18

beryllium - atomic # 4 calcium - atomic # 20 - mass # 9.012 = 9 - mass # 40.08 = 40 # p+ = 4 20# p+ = # e- = 4 20# e- = #n° = 9 – 4 = 5 40#n° – 20 = = 20

2 e- 8e- 20 e- 2 e- 8 e- 4 e- 2 e- 2 e- p+ = 4 p+ = 20 n° = 5 n° = 20 H. Octet Rule  atoms tend to be stable when the outer energy level is full of electrons  the octet rule states that atoms bond in a way to have a full valence energy level (there are exceptions)  atoms will either share electrons, or gain or lose electrons in order to satisfy this octet rule  compounds are formed when two or more different elements bond together either by sharing or transferring electrons

I. Ions  ions are atoms or groups of atoms that have a net charge  number of p+ and e are not equal  most atoms try to achieve the electron configuration of a noble gas  isoelectronic means having the same number of e as another atom or ion eg) fluorine gains an electron to be isoelectronic with neon eg) potassium loses an electron to be isoelectronic with argon

Cations  positively charged ions  lost electrons to obtain a stable electron configuration (full valence level)  METALS form cations eg) sodium loses one e to completely empty the last energy level Na+ has 10 e and 11 p+  charge on a metal ion is the same as the group number for groups 1,2,3,13,14 (ignore the 1 in front of 13 and 14)

Anions  negatively charged ions  gained electrons to obtain a stable electron configuration  NON-METALS form anions  have “ide” ending eg) oxygen gains two e to completely fill the last energy level O2 has 10 e and 8 p+ and is called the oxide ion  charge on a non-metal is 18  group number (not Roman Numerals) EELR’s for Ions number of p+ = atomic number number of e = number of p+ – charge number of n0 = atomic mass – atomic number (# of p+)

Examples

nitride ion - atomic # 7 -7 mass # 14.007 = 14 # p+ = 7-(-3) = 10 # e = 14 – 7 = 7 #n° =

8 e- 10 e- 2 e- p+ = 7 n° = 7

calcium ion - atomic # 20 - mass # 40.08 = 40 # p+ = 20 # e- = 20 – (+2) = 18 #n° = 40 – 20 = 20

8 e- 8 e- 18 e- 2 e- p+ = 20 n° = 20 Draw the EELR for the following: 1. sulphide ion 2. aluminium ion 3. chloride ion 4. magnesium ion sulphide ion - atomic # 16 aluminium ion - atomic # 13 16 - mass # 32.07 = 32 13 - mass # 26.98 = 27 # p+ = 16 – (2) = 18 # p+ = 13 – (+3) = 10 # e- = 32 – 16 = 16 # e- = 27 – 13 = 14 #n° = #n° = 8 e- 8 e- 18 e- 8 e- 10 e- 2 e- 2 e- p+ = 16 p+ = 13 n° = n° = 14 16

chloride ion - atomic # 417 magnesium ion - atomic # 12 17 - mass # 35.45 = 35 12 - mass # 24.31 = 24 # p+ = 17 – (1) = 18 # p+ = 12 – (+2) = 10 # e- = 35 – 17 = 18 # e- = 24 – 12 = 12 #n° = #n° = 8 e- 8 e- 18 e- 8 e- 10 e- 2 e- 2 e- p+ = 17 p+ = 12 n° = n° = 12 18

Your Assignment: pg 2 J. Elements  metallic elements exist as single atoms (monatomic)  chemical formula is simply the symbol followed by the state at room temperature eg) sodium Na(s) mercury Hg(ℓ) copper Cu(s)  non-metals (not including noble gases) do not exist as single atoms and are called molecular elements (diatomic, polyatomic)  chemical formula is the symbol with the subscript and the state at room temperature  memorize the subscripts (flagpole)!!!

H 2 N O F 2 2 2 P S Cl 4 8 2 Br 2 I 2

Summary: Monatomic C(s), noble gases, all metals

Diatomic H2(g), N2(g), O2(g), F2(g), Cl2(g), Br2(g), I2(s)

Polyatomic P4(s), S8(s)

K. Molecular Compounds  molecular compounds are formed when two or more nonmetals bond together  bonded by covalent bonds which is the force of attraction between atoms that are sharing electrons  properties: 1. do not conduct electricity when dissolved in water 2. dissolve in water to form either a neutral molecular solution or an acidic solution 3. solids, liquids or gases at room temperature Naming  give the atom name for the first element (with the prefix if there is more than one) then give the name for the second element with “ide” ending and include the prefix  Note: if the first element is hydrogen, do not put a prefix (these are acids!)

Prefixes 1 = mono 4 = tetra 7 = hepta 2 = di 5 = penta 8 = octa 3 = tri 6 = hexa 9 = nona 10 = deca

Try These

1. CO(g) carbon monoxide 2. CO2(g) carbon dioxide 3. P4O10(s) tetraphosphorus decaoxide 4. BrH7(s) bromine heptahydride

Writing Formulas  simply write each symbol followed by the subscript (from prefix)

Try These

1. oxygen dibromide OBr2 2. diphosphorus pentasulphide P2S5 3. carbon tetraiodide CI4 4. phosphorus pentachloride PCl5

 some molecular compounds have classical names…memorize them!!! O3 = ozone NH3 = ammonia H2O = water H2O2 = hydrogen peroxide CH4 = methane CH3OH = methanol C2H6 = ethane C2H5OH = ethanol C6H12O6 = glucose C12H22O11 = sucrose H2S = hydrogen sulphide HCl, HF, HBr, HI = hydrogen halides (usually acids)

Your Assignment: pg 3 L. Ionic Compounds  ionic compounds are formed when electrons are transferred, allowing oppositely charged ions to bond together  ionic bond is the force of attraction between oppositely charged ions  properties: 1. conduct electricity when dissolved in water 2. separate into ions when dissolved in water 3. crystalline solids at room temperature

1. Monovalent Ionic Compounds  monovalent means there is one charge on the metal eg) Na+, Ca2+  metal + nonmetal

Naming  give the name for each ion…metals are normal, non-metals have “ide” ending *** never use capital letters in naming!! eg) NaF sodium fluoride

Na2S sodium sulphide the 2 means that two sodium ions are bonded with one sulphide ion… this doesn’t matter for naming Try These: Name the following: 1. LiF lithium fluoride 2. KCl potassium chloride 3. BeS beryllium sulphide

4. Rb3P rubidium phosphide 5. MgF2 magnesium fluoride 6. Na2O sodium oxide 7. CsBr cesium bromide

8. BaCl2 barium chloride

Writing Formulas  look up the symbol for each ion and write them listing the metal ion first  balance the charges using subscript numbers ***in ionic compounds, the total positive charges must equal the total negative charges…the net charge is zero eg) sodium oxide Na2O

1+ 2 1+  2 = 2+ 2  1 = 2

calcium phosphide Ca3P2

2+ 3 2+  3 = 6+ 3  2 = 6

Try these:

1. magnesium chloride MgCl2 2. calcium chloride CaCl2 3. zinc sulphide ZnS 4. silver sulphide Ag2S 5. germanium oxide GeO2 6. calcium arsenide Ca3As2 7. magnesium nitride Mg3N2

2. Multivalent Ionic Compounds  metal ions that have more than one possible charge eg) Cu2+, Cu+, Fe3+, Fe2+  transition metal + nonmetal  the first charge listed is the most common

Naming  same rules as before  the difference…you must include brackets containing the charge of the metal ion in Roman Numerals (I,II,III,IV,V,VI,VII)  figure out the charge on the metal by using how many negative charges there are in the nonmetal ions eg) CuI copper (I) iodide

TiBr4 titanium (IV) bromide Ti3P4 titanium (IV) phosphide Try These: 1. AuBr gold (I) bromide

2. CrCl2 chromium (II) chloride 3. Co2O3 cobalt (III) oxide 4. VS2 vanadium (IV) sulphide 5. PuN2 plutonium (VI) nitride

Writing Formulas  same rules  the charge on the metal is given to you in the brackets eg) iron (II) oxide FeO

tin (II) chloride SnCl2 chromium (III) sulphide Cr2S3

Try these:

1. chromium (II) sulphide CrS 2. nickel (III) chloride NiCl3 3. vanadium (IV) phosphide V3P4 4. gold (III) iodide AuI3

Your Assignment: pg 4

3. Mixed Ionic Compounds  metal ion + polyatomic ion (complex ion) 3 2  eg) PO4 , SO4 , HCO3 etc. + ***NH4 (ammonium ion) is the only positive complex ion…you’ll see it in the place of a metal

Naming  give the name for the first ion then give the name for the complex ion

Try These:

1. KIO3 potassium iodate 2. NaCH3COO sodium acetate 3. MgSO3 magnesium sulphite 4. NH4NO3 ammonium nitrate 5. Ca3(PO4)2 calcium phosphate Writing Formulas  same as before…look up the symbol for each ion then balance the charges using subscripts  if you must multiply one of the complex ions, put brackets around it first then write the subscript

Try These:

1. aluminum phosphate AlPO4 2. aluminum chlorate Al(ClO3)3 3. calcium sulphite CaSO3 4. scandium acetate Sc(CH3COO)3 5. ammonium sulphate (NH4)2SO4

4. Hydrated Ionic Compounds  ionic compounds containing water in their structure

 water is represented by “xH2O” in the formula where x is the number of water molecules

Naming  give the name for the first part of the compound using our previous rules

 name the “xH2O” part as prefix + “hydrate” (same prefixes as molecular compounds)

Try These:

1. NaF3H2O sodium fluoride trihydrate

2. FeSO46H2O iron (II) sulphate hexahydrate

3. GaNH2O gallium nitride monohydrate

Writing Formulas  same as before…look up the symbol for each ion then balance the charges using subscripts

 for the hydrate part…add “xH2O” where x is the number given in the prefix Try These:

1. iron (III) nitrate nonahydrate Fe(NO3)39H2O

2. sodium chlorate tetrahydrate NaClO34H2O

3. nickel (II) sulphite heptahydrate NiSO37H2O

Summary: Ionic vs. Molecular Ionic Molecular  metal + non-metal (or  all non-metals polyatomic ions)  no prefixes (except hydrates)  prefixes  solids  solids, liquids or gases  solutions conduct  solutions do not conduct  solutions are basic or neutral  solutions are acidic or neutral

5. Acids and Bases  matter can be subdivided into three groups based on its properties: 1. acids 2. bases 3. neutral substances  let’s look at the properties of acids and bases:

Acids Bases

 usually soluble in H2O  usually soluble in H2O  conduct electricity  conduct electricity  neutralize bases  neutralize acids  taste sour  taste bitter  react with metals to produce  feel slippery H2(g) Indicators Indicators  litmus – pink  litmus – blue  bromothymol blue – yellow  bromothymol blue – blue  phenolphthalein – colourless  phenolphthalein – bright pink eg) vinegar, lemon juice eg) ammonia, Tums  the pH scale was devised to indicate how acidic or basic a substance is

stomach acid coffee antacid KOH strong acids 0 1 5 7 11 14strong bases acids bases Neutral Naming and Writing Formulas for Bases  use rules for ionic compounds  many of them contain the hydroxide ion (OH) eg) NaOH sodium hydroxide KOH potassium hydroxide

Ba(OH)2 barium hydroxide NaHCO3 sodium hydrogen carbonate calcium hydroxide Ca(OH)2 calcium carbonate CaCO3

Naming Acids  acids always contain hydrogen (usually as the first element)  acids are always aqueous (dissolved in water) (aq) eg) HCl(aq), H2SO4(aq), HNO3(s) not an acid

Rules hydrogen ______ide becomes hydro______ic acid hydrogen ______ate becomes ______ic acid hydrogen ______ite becomes ______ous acid

Try These: Write the acid name for the following: 1. HF(aq)  hydrogen fluoride  hydrofluoric acid

2. H2SO3(aq)  hydrogen sulphite  sulphurous acid

3. H3BO3(aq)  hydrogen borate  boric acid 4. HCl(g)  hydrogen chloride  not an acid!!!!!!!! Writing Formulas  use the naming acids rules in the opposite direction  come up with the “ionic” name then write the formula, balance the charges and add “(aq)” to the end

Try These: Write the ionic name and formula for the following:

1. hydrosulphuric acid hydrogen sulphide H2S(aq) 2. carbonic acid hydrogen carbonate H2CO3(aq) 3. chlorous acid hydrogen chlorite HClO2(aq)

***Note the special cases of sulphur and phosphorus…they add “ur” and “or” respectively to the name

Your Assignment: pgs 7-8

M. States and Solubility  acids – always (aq)  elements – can be (s), (l) or (g)…look up state on periodic table  molecular compounds – can be (s), (l), or (g)…the question usually tells you (or use common sense!)  ionic compounds – either (s) or (aq)…look up on the solubility chart

Examples

1. LiCl (aq) 4. Ba(OH)2(aq) 2. AgCl(s) 5. BaSO4(s) 3. NaNO3(aq) 6. K2S(aq )

N. Chemical Reactions  chemical reactions can cause both physical and chemical changes and always involve the formation of a new substance

Evidence 1. temperature change 2. colour change 3. solid (precipitate) produced 4. gas produced Reactants Products balancing

1 Zn + 2 HCl  1 ZnCl + 1 H (s) (aq) 2(aq) 2(g)

states states

 energy changes that occur with chemical reactions can be: 1. endothermic = energy is absorbed (enters) eg) cooking food reactants + energy  products

2. exothermic = energy is released (exits) eg) combustion reactants  products + energy

O. Law of Conservation of Matter  Law of Conservation of Matter states that matter cannot be created or destroyed, it only changes forms  mass of reactants = mass of products

Counting Practice! How many of each element are in the following compounds?

1. NaCl 4. Ba(OH)2 7. 2 CaCl2 2. BaBr2 5. NH4CH3COO 8. 8 PbI2 3. (NH4)3P 6. 3 (NH4)2S 9. 4 Zn(CH3COO)2

 when chemicals react they follow the Law of Conservation of Matter  there must be equal numbers of each element on both sides of the reaction  coefficients are used to increase the number of atoms in a compound (balancing)

Examples

1. 2 Mg(s) + 1 O2(g)  2 MgO(s)

 Mg 1 x 2 = 2 Mg 1 2 O 2 O 1 x 2 = 2  2. 2 H2O(l)  2 H2(g) + 1 O2(g)

 H 2 x 2 = 2 H 2 O 1 2 O 2 x 2 = 2 

3. 1 Cu(s) + 2 AgNO3(aq)  2 Ag(s) + 1 Cu(NO3)2(g)

Cu 1 Cu 1 Ag 1 2 Ag 1 x 2 = 2   NO 1 x 2 = 2 NO 2 3 3

4. 1 Cl2(g) + 2 NaBr(aq)  1 Br2(l) + 2 NaCl(aq)

Cl 2 Cl 1 2 Na 1 2 Na 1 x 2 = 2   Br 1 x 2 = 2 Br 2

5. 2 KI(aq) + 1 Pb(NO3)2(aq)  1 PbI2(s) + 2 KNO3(aq)

 K 1 x 2 = 2 K 1 2 I 1 2 I 2 NO 2 NO 2 x 2 = 2  3 3 Pb 1 Pb 1

6. 1 CH4(g) + 2 O2(g)  1 CO2(g) + 2 H2O(g)

C 1 C 1 H 4 H 2 x 2 = 4   O 2 x 2 = 4 O 2 + 1 = 3 2 + 2 = 4

Your Assignment: pg 10 P. Identifying Chemical Reactions 1. Hydrocarbon Combustion  a compound containing hydrogen and carbon (a hydrocarbon) burns/combusts in the presence of O2(g)

 always forms CO2(g) and H2O(g)

C?H? + O2(g)  CO2(g) + H2O(g)

eg) 1 CH4(g) + 2 O2(g)  1 CO2(g) + 2 H2O(g)

2. Composition, Combination or Formation  will always have more reactants than products  are usually exothermic i. Combination of Elements  elements combine to form a compound element + element  compound

eg) 2 Mg(s) + 1 O2(g)  2 MgO(s)

ii. Combination to Produce an Acid (AP)  combination of a non-metal oxide (acidic oxide) with water non-metal oxide + water  acid

eg) 1 CO2(s) + 1 H2O(ℓ)  1 H2CO3(aq) iii. Combination to Produce a Base (AP)  combination of a metal oxide (basic oxide) with water metal oxide + water  base

eg) 1 Li2O(s) + 1 H2O(ℓ)  2 LiOH(aq) iv. Combination to Produce a Salt (AP)  combination of a metal oxide (basic oxide) with a non-metal oxide (acidic oxide) metal oxide + non-metal oxide  a salt

eg) 1 MgO(s) + 1 CO2(g)  1 MgCO3(s) 3. Decomposition  will always have more products than reactants  are usually endothermic i. Decomposition to Produce Elements  compound decomposes into its elements compound  element + element

eg) 2 H2O(l)  2 H2(g) + 1 O2(g)

ii. Decomposition of Acids (AP)  acid decomposes into non-metal oxide (acidic oxide) and water acid  non-metal oxide + water

eg) 1 H2SO4(aq)  1 SO3(g) + 1 H2O(l)

iii. Decomposition of Bases (AP)  base decomposes into metal oxide (basic oxide) and water base  metal oxide + water

eg) 1 Ca(OH)2(s)  1 CaO(s) + 1 H2O(l)

iv. Decomposition of Salts (AP)  a salt decomposes into non-metal oxide (acidic oxide) and metal oxide (basic oxide) a salt  non-metal oxide + metal oxide

eg) 1 Fe2(SO4)3(s)  3 SO3(g) + 1 Fe2O3(s)

4. Single Replacement (Displacement)  an element reacts with an ionic compound to form a different element and a different ionic compound element + compound  element + compound

eg) Cu(s) + 2 AgNO3(aq)  2 Ag(s) + Cu(NO3)2(g)

Cl2(g) + 2 NaBr(aq)  Br2(l) + 2 NaCl(aq) 5. Double Replacement  two ionic compounds react to form two different ionic compounds compound + compound  compound + compound i. Precipitation Reactions  one of the products formed is insoluble compound + compound  insoluble compound + compound

eg) KI(aq) + Pb(NO3)2(aq)  PbI2(s) + KNO3(aq)

ii. Gas Formation (AP)  sometimes one of the products in a double replacement undergoes spontaneous decomposition, forming a gas  you must write down the products of this decomposition  memorize these!

H2CO3(aq)  H2O(ℓ) + CO2(g)

NH4OH(aq)  H2O(ℓ) + NH3(g)

H2SO3(aq)  H2O(ℓ) + SO2(g)

eg) CaCO3(s) + HNO3(aq)  Ca(NO3)2(aq) + H2O(ℓ) + CO2(g) iii. Neutralization Reactions  an acid reacts with a base to produce water and a salt acid + base  water + a salt

eg) 1 HCl(aq) + 1 NaOH(aq)  1 H2O(l) + 1 NaCl(aq)

1 ZnO(s) + 2 HNO3(aq)  1 H2O(l) + 1 Zn(NO3)2(aq)

2 LiOH(aq) + 1 CO2(g)  1 H2O(l) + 1 Li2CO3(aq)

Your Assignment: pgs 11-13 Q. Significant Digits  any digit from 1 – 9 is significant  trailing zeros are significant eg) 6.3800, 12 000  “sandwich” zeros are significant eg) 2.04, 1005.002  leading zeros are not significant eg) 0.0032  counted objects and constants are not included in sig digs

 / : multiply of divide then round answer to the lowest number of sig digs

 +/ : add or subtract then round answer to the lowest number of decimal places

R. The Mole  the mass of a single atom is so small that we cannot easily measure it  the mole, n, is a measurement that is used so that we can actually measure the mass of elements and compounds…IT IS JUST A NUMBER!!!

1. Avogadro’s Number 23  1 mole = Avogadro’s number (NA) = 6.02 x 10 atoms, molecules etc. that’s 602 000 000 000 000 000 000 000 atoms, molecules etc!

you can use Avogadro’s number to calculate the number of moles in a substance if you know the number of molecules

n = # atoms

Example A diamond contains 5.0 x 1025 atoms of carbon. How many moles of carbon are in this diamond?

n = 5.0  10 25 atoms 6.02  1023 atoms/mol = 83.056….mol = 83 mol

Your Assignment: pg 15 #1-3 2. Atomic Molar Mass  12C, carbon 12, has been assigned by chemists (1961) as the standard for measuring atomic masses  12C is assigned a mass of exactly 12 atomic mass units (amu)  comparing masses of atoms involves the use of a mass spectrometer  an electron beam and magnetic field are used to deflect ions… amount of deflection indicates the mass (higher mass = lower deflection)

 once you know the mass of each isotope, they can then be averaged (weighted) to come up with the average atomic mass for the element  the atomic masses given on the periodic table are an average of all the naturally occurring isotopes of each element

Example 1 It is known that natural carbon is composed of 98.89% 12C at 12 amu and 1.11% 13C at 13.003355 amu. The amount of 14C is so small it is not used in these calculations. Calculate the average atomic mass of natural carbon.

atomic mass = 98.89% of 12 amu + 1.11% of 13.003355 amu = (0.9889)(12 amu) + (0.0111)(13. 003355 amu) = 12.01113724 amu = 12.0 amu

Example 2 When a sample of natural copper is vaporized and injected into a mass spectrometer, the sample is found to be 69.09% 63Cu with a mass of 62.93 amu and 30.91% 65Cu with a mass of 64.93 amu. Calculate the average atomic mass of copper.

atomic mass = 69.09% of 62.93 amu + 30.91% of 64.93 amu = (0.6909)(62.93 amu) + (0.3091)(64.93 amu) = 63.5482 amu = 63.55 amu

Your Assignment: pg 15 #4-5  molar mass is the average atomic mass of one mole of a substance  measured in g/mol  the average atomic molar masses are given on the periodic table

Examples: Calculate the molar mass of the following substances:

1. Na(s) 1 x 22.99 g/mol = 22.99 g/mol

2. O2(g) 2 x 16.00 g/mol = 32.00 g/mol

3. Al2O3(s) 2 x 26.98 g/mol = 53.96 g/mol 3 x 16.00 g/mol = 48.00 g/mol 101.96 g/mol

4. Ca(OH)2(s) 1 x 40.08 g/mol = 40.08 g/mol 2 x 16.00 g/mol = 32.00 g/mol 2 x 1.0079 g/mol = 2.0158 g/mol 74.10 g/mol

3. Mole Calculations  now we can use number of moles and molar mass in a formula:

n = m M m = nM where: n = number of moles in mol m = mass in g M = molar mass in g/mol

Example 1: How many moles are in 200 g of table salt (NaCl)? m = 200 g n = m M = 1 x 22.99 g/mol = 22.99 g/mol M 1 x 35.453 g/mol = 35.453 g/mol = 200 g 58.443 g/mol 58.443 g/mol = 3.42 mol Example 2:

How many grams are in 62.9 mol of lead (II) nitrate? Pb(NO3)2 n = 62.9 mol m = nM M = 1 x 207.19 g/mol = 207.2 g/mol = (62.9 mol)(331.2 g/mol) 2 x 14.01 g/mol = 28.02 g/mol = 2.08  104 g 6 x 16.00 g/mol = 96.00 g/mol 331.2 g/mol

Your Assignment: pg 15-17 #4-14

S. Gravimetric Stoichiometry  gravimetric = mass measurements

 when you balance a reaction, you are balancing the moles of each reactant and product…this gives you the mole ratio

eg) 2 Mg(s) + 1 O2(g)  2 MgO(s) mole ratio is 2:1:2

 certain assumptions must be made: 1. the reaction is spontaneous 2. the reaction is fast 3. the reaction goes to completion 4. the reaction is stoichiometric ie) whole number ratio of moles of reactants and products

Steps 1. Write a balanced equation including the states. Write the information given. 2. Find the moles of the given species using n=m/M. 3. Find the moles of the wanted species using mole ratio (wanted/given). 4. Calculate mass of the wanted species using m=nM.

Example 1 Calculate the moles of aluminum formed when 2.50 moles of aluminum oxide decomposes into aluminum and oxygen.

2 Al2O3(s)  3 O2(g) + 4 Al(s) x mol n = 2.50 mol n = 2.50 mol  4 2 = 5.00 mol Your Assignment: pg 18

Example 2 Calculate the mass of sulfur needed to react with 0.500 moles of zinc.

8 Zn(s) + 1 S8(s)  8 ZnS(s) x g n = 0.500 mol n = 0.500 mol  1 8 = 0.0625 mol M = 256.56 g/mol m = nM = (0.0625 mol)(256.56 g/mol) = 16.035 g = 16.0 g

Example 3 Iron is produced by the reaction of iron(III) oxide with carbon monoxide to produce iron and carbon dioxide. What mass of iron(III) oxide is required to produce 1000 g of iron?

Fe2O3(s) + 3CO(g)  2Fe(s) + 3CO2(g) x g M = 159.70 g/mol m = 1000 g n = 17.90… x 1/2 M = 55.85 g/mol = 8.95… mol n = m/M m = nM = 1000 g/55.85g/mol = (8.95… mol)(159.70 g/mol) = 17.90… mol = 1429.72… g = 1430 g

Your Assignment: pgs 20-21 Chemistry Review: pgs 22-24