CHEM110W2 Ms Janine Kasavel Rm:

Quantitative Chemistry matter units significant figures atomic structure, isotopes, periodic table basic nomenclature (ions, molecular &inorganic compounds) stoichiometry and balancing equations by inspection moles and Avogadro’s number empirical and molecular formulae limiting reagents 2

MATTER “matter is anything that has mass and takes up space” Pure Substance -ELEMENT: can’t be decomposed into simpler substances -COMPOUND: composed of 2/> different elements Mixture -HETEROGENEOUS: visibly different composition, properties or appearance -HOMOGENEOUS: visibly uniform composition, properties & appearance throughout 3

Physical properties of matter -measured without changing the identity or composition of the substance Chemical properties of matter -describe the way a substance may change or react to form other substances 4

UNITS Système International (SI) MASS kilogram kg LENGTH metre m TIME second s TEMPERATURE Kelvin K G giga 10 9 M mega 10 6 k kilo 10 3 d deci c centi m milli µ micro n nano p pico SCIENTIFIC NOTATION

SIGNIFICANT FIGURES 1)any figure that is not zero is significant. 2)zeroes between non-zero figures are significant. 3) exact (“counting”) numbers by definition have an ¥ number of s.f., so physical constants defined to be exact numbers do so also. 4) leading zeroes (to the left of the first non-zero figure) are not significant. 5)trailing zeroes (to the right of the last non-zero figure) are significant only if the number has a d.p. 6)in measurements without a d.p., the number of s.f. is ambiguous. 6

Using Significant Figures in Calculations multiplication/division Number of s.f. in final answer is the same as the LEAST of numbers of s.f. in each of original measurements. addition/subtraction Number of d.p. in final answer is the same as the LEAST of numbers of d.p. in each of original measurements. 7

DENSITY ρ = mass/ volume ρ: gcm -3 mass: g volume: cm 3 8

PRACTICE EXAMPLE 9 A nugget of gold with a mass of 521 g is added to 50.0 mL of water. The water level rises to a volume of 77.0 mL. What is the density of the gold?

ATOMIC STRUCTURE 10 PROTONS, NEUTRONS in the nucleus surrounded by orbiting ELECTRONS Early Atomic Theory (Dalton 1803 – 1807) Cathode Rays & Particles (Thomson, 1897) Electron Charge & Mass (Millikan, 1909) Nuclear Atom (Rutherford, 1910) Modern Atomic Structure (Rutherford, 1919)

11 ChargeMass Actual/ Coulombs Relative Actual/ g Relative/ u Proton1.602 x x Electron1.602 x x Neutron x

12 A: mass number = no. protons + no. neutrons Z: atomic number = no. protons /electrons A Z E

Isotopes Atoms of the same element with different mass numbers due to: different numbers of neutrons 13

Average atomic mass 14 AAM: average atomic mass IM: isotopic mass

PRACTICE EXAMPLE 15 Naturally occurring Mg has three isotopes: 24 Mg (78.90 %) u 25 Mg (10.00 % ) u 26 Mg (11.10 %) u AAM=?

16

17 IONS If electrons are added to or removed from a neutral atom, an ion is formed. When an atom or molecule loses electrons it becomes positively charged  CATION (E + ) 11 p + 11 e - 11 p + 10 e - Na atom Na + ion L.Pillay 2010

18 When an atom or molecule gains electrons it becomes negatively charged  ANION (E - ). Generally, metal atoms tend to lose electrons (forms cations) and non-metal atoms gain electrons (forms anions). 17 p + 18 e - 17 p + 17 e - Cl atomCl - ion L.Pillay 2010

19 COMMON CATIONS

20 COMMON ANIONS

Ionic compounds Composed of nonmetal and metal Cations and anions attract each other to form a neutral compound NAMES: Name of metal (cation) written first If metal has more than one common charge, write the charge in roman numerals in brackets Name of nonmetal (anion) written next with –ide ending FORMULAE: compounds are electrically neutral, the formula of a compound can easily be constructed simply by: -writing value of cation charge as subscript on anion -writing value of anion charge as subscript on cation 21

PRACTICE EXAMPLE 22 NaCl K 2 SO 4 Ba(OH) 2 cobalt(II) nitrate silver sulfide ferric chloride

Oxyanion ClO 4 - perchlorate ion (one more O atom than chlorate) ClO 3 - chlorate ion (one more O atom than chlorite) ClO 2 - chlorite ion (one more O atom than hypochlorite) ClO - hypochlorite ion Acids -acids containing anions whose names end in -ide are named by changing the -ide ending to -ic, adding the prefix hydro- to this anion name, and then following with the word acid - acids containing anions whose names end in -ate/-ite are named by changing the -ate ending to -ic or the -ite ending to -ous and then adding the word acid 23

PRACTICE EXAMPLE 24 AnionCorresponding acid Cl - S 2- ClO 4 - ClO 3 - ClO 2 - ClO -

MOLECULAR COMPOUNDS Generally composed only of nonmetals Diatomic species includes O 2 N 2, F 2, Br 2, I 2 NAMING: name of element furthest left on periodic table generally written first both elements in same group on periodic table, element with higher Z written first name of 2nd element given the ending –ide Greek prefixes used to indicate number of atoms of each element Greek prefixes: mono, di, tri, tetra, penta, hexa, hepta, octa, nona, deca

PRACTICE EXAMPLE 26 SO 2 PCl 5 N 2 O 3 NF 3 P 4 S 10 silicon tetrabromide

STOICHIOMETRY “quantities of substances consumed and produced in chemical reactions” Atoms are neither created or destroyed in a chemical reaction. A chemical equation must have equal numbers of atoms of each element on each side of the arrow. The molecular composition of certain ions must remain the same on each side of the arrow. 27

PRACTICE EXAMPLE 28 C 2 H 6 + O 2 → CO 2 + H 2 O Al + HCl → AlCl 3 + H 2

MOLE & AVOGADRO’S NUMBER Number of atoms/molecules/ions represented as mole amounts Avogadro’s number: N A = X mol 12 C atoms = X C atoms 1 mol H 2 O molecules = X H 2 O molecules 1 mol NO 3 - ions = X NO 3 - ions 29

Molar mass “Mass in grams of one mole of a substance” Related to mole amount of a substance by the equation: 30 n: number of moles (in mol) m: mass (in grams) MM: molar mass (in grams per mole) m n MM

PRACTICE EXAMPLE 31 How many oxygen atoms are in 1.50 mol of sodium carbonate?

EMPIRICAL AND MOLECULAR FORMULA “Ratio of atoms of each element in a compound” 32 Mass % elements Grams of each element Moles of each element Empirical formula Assume Use molar Calculate 100g mass mole ratio sample

PRACTICE EXAMPLE 33 Determine the empirical formula of a compound with 10.4% C, 27.8% S and 61.8% Cl.

QUIZ 2 Wed 17 August during the tut Stoichiometry (slide 28) until end quantitative chemistry TUTS Check chem foyer for new tut group lists after 1pm on Monday 15/08 34

PRACTICE EXAMPLE 35 Eucalyptol has an empirical formula of C 10 H 18 O. The experimentally determined molecular mass of this substance is 152 u. What is its molecular formula?

36

LIMITING REACTANTS In chemical reactions one reactant (the LIMITING reactant (LR)) is used up first The reaction stops once LR used up leaving the excess reactants as leftovers The LR limits the amount of product formed 37

1.Balance equation 2.Determine the molar masses 3.Calculate the number of moles for each reactant and divide by stoichiometric coefficients 4.Smallest mole amount =LR 38

PRACTICE EXAMPLE Determine (i) which reactant is the limiting reactant and (ii) how much excess reactant is leftover in the following reaction: 3NH 4 NO 3 + Na 3 PO 4 → (NH 4 ) 3 PO 4 + 3NaNO g 50.0 g 39

Theoretical yield Quantity of product that is calculated to form, based on LR Actual yield is the mass of product obtained when the experiment is carried out Percent yield relates actual yield to theoretical yield % Yield = (actual yield/theoretical yield) x 100% 40

PRACTICE EXAMPLE Determine the theoretical yield and the % yield of NaNO 3 if 15.0 g of sodium nitrate is formed when the reaction is carried out. 3NH 4 NO 3 + Na 3 PO 4 → (NH 4 ) 3 PO 4 + 3NaNO g 50.0 g 41

QUANTITATIVE CHEMISTRY REVISION CHAPTER 1 and 2 42