Unit 1 The Structure of Matter

Units of Measurement SI Units Physical QuantityUnit NameSymbol lengthmeterm masskilogramkg timeseconds electric currentampereA temperatureKelvinK amount of substancemolemol luminous intensitycandelacd

Metric System Conversions

Generally require temperatures are in Kelvins T (K) = t (˚C) Generally require temperatures are in Kelvins T (K) = t (˚C) Temperature Conversions Celsius 100 ˚C 0 ˚C 100˚C Kelvin 373 K 273 K 100 K Fahrenheit 32 ˚F 212 ˚F 180˚F

Three targets with three arrows each to shoot. Can you hit the bull's-eye? Both accurate and precise Precise but not accurate Neither accurate nor precise How do they compare? Can you define accuracy and precision?

All nonzero numbers and zeros between nonzero digits are significant (four siggy figgies in each) Zeros to the left of the first nonzero number are not significant (one siggy figgy and two siggy figgies) Zeros at the end of a number to the right of the decimal are significant (five siggy figgies and four siggy figgies) Zeros at the end of number greater than one are not significant unless their significance is indicated by the presence of a decimal point (one siggy figgy and three siggy figgies) Significant Figures

The coefficients of a balanced equation and numbers obtained by counting objects are infinitely significant Using Significant Figures in Computations When multiplying or dividing – the result should have the same number of significant figures as the number in the calculation with the smallest number of significant figures.

When adding or subtracting – the result should have the same number of decimal places as the number in the calculation with the smallest number of decimal places.

Why Are Significant Figures Significant?

More Practice

Calculations & Conversion Practice Determine the length in kilometers if a x m automobile race. Watch for significant figures and be sure to include units on your answer.

Calculations & Conversion Practice Calculate the mass of 1.0L of benzene if it has a density of 0.879g/mL

Calculations & Conversion Practice If the volume of an object is reported as 5.0 cm 3, what is the volume in m 3 ?

Calculations & Conversion Practice If the volume of an object is reported as 4.5m 3, what is the volume in nm 3 ?

Classification of Matter

2.6 – Molecules and Molecular Compounds (Covalent Compounds) Two or more atoms tightly bound together Bond by a covalent bond – the sharing of electrons Usually nonmetals bonded to other nonmetals Elements found in nature in molecular form – N 2, O 2, F 2, Cl 2, Br 2, I 2, H 2 Previous list is called diatomic elements Molecular formulas – indicate actual numbers of and types of atoms in a molecule Term “molecule” refers only to covalently bonded substances. Empirical formulas – smallest possible whole number subscripts Use the Greek prefixes to name binary covalent compounds

Prefixes Used in Naming Molecular Compounds NumberPrefix 1mono- 2di- 3tri- 4tetra- 5penta- 6hexa- 7hepta- 8octa- Example Give the chemical formula for the following: silicon tetrachloride disulfur dichloride

2.7 – Ions and Ionic Compounds Ion – formed when electrons are added or removed from an atom Cation – ion with a positive charge – typically metals Anion – ion with a negative charge – typically non-metals Polyatomic ions – atoms joined as a molecule but have a net positive or negative charge Use the periodic table to predict charges Formed by the transfer of electrons Compounds formed when cations and anions are attracted to each other Typically formed by metals and nonmetals

Naming Ionic Compounds Cations – metals – have the same name as the metal Must use roman numerals if dealing with a metal that can have various charges – usually the transition metals One atom anions – nonmetals – ending changes to –ide Must Memorize the polyatomic chart below!

Dealing With Various Amounts of Oxygen Example The formula for bromate ion is BrO 3 -. What is the formula for hypobromite ion? Example Name the following: NH 4 Br, Cr 2 O 3, Co(NO 3 ) 2

Example Give the chemical formula for the following: magnesium sulfate silver sulfide lead(II)nitrate

Naming acids I –ate it and it was –ic I caught –ite –ous Binary Acids Hydro ……. -ic Example Give the formulas for the following: 1.hydrobromic acid 2.carbonic acid 3.Sulfurous acid

Stoichiometry – the math of chemistry Law of Conservation of Mass – the total mass of all substances present after a chemical reaction is the same as the total mass before the reaction. Atomic Theory – atoms are neither created nor destroyed 3.1 – Chemical Equations Fe (s) + O 2 (g)  Fe 2 O 3 (s) C 2 H 4 (g) + O 2 (g)  CO 2 (g) + H 2 O(g) Al (s) + HCl(aq)  AlCl 3 (aq) + H 2 (g)

3-4 – The Mole 1 mol element = 6.02 x atoms 1mol molecule = 6.02 x molecules 1 mol ions = 6.02 x ions How many oxygen atoms are in 0.25 mol of Ca(NO 3 ) 2 ? How many oxygen atoms are in 1.50 mol of sodium carbonate?

Molar mass (a.k.a gram formula mass) The mass of one mole of any substance Units are g/mol Calculated from periodic table Examples Calculate the molar mass of Ca(NO 3 ) 2 How many moles of sodium bicarbonate are there in 508g of sodium bicarbonate? What is the mass in grams of 6.33 mol of NaHCO 3 ? What is the mass in grams of 3.0 x mol of sulfuric acid? How many nitric acid molecules are in 4.20g of HNO 3 ? How many oxygen atoms are there?

3.5 – Empirical and Molecular Formulas Example A g sample of methyl benzoate, a compound used in the manufacturing of perfumes, is found to contain 3.758g of carbon, 0.316g of hydrogen and 1.251g of oxygen. What is the empirical formula of this substance?

Example Ethylene glycol, the substance uses in automobile antifreeze, is composed of 38.7% carbon, 9.7% hydrogen and 51.6% oxygen by mass. Its molar mass is 62.1g/mol. What is the empirical formula? What is the molecular formula?

Example Caprice acid, which is responsible for the foul odor of dirty socks, is composed of C, H and O atoms. Combustion of g sample of this compound produces 0.512g of CO 2 and 0.209g of H 2 O. What is the empirical formula of caproic acid? The molar mass is 116g/mol. What is the molecular formula?

3.6 – Quantitative Information Coefficients of balanced equations Relative number of molecules Relative number of moles Use in mole ratio to convert between substances in a stoichiometry calculation Example The decomposition of KClO 3, is commonly used to prepare small amounts of O 2 in the laboratory. 2KClO 3 (s)  2 KCl (s) + 3O 2 (g) How many grams of O 2 can be prepared from 4.50g of KClO 3 ?

Example Propane is a common fuel used in cooking and heating. What mass of O 2 is consumed in the combustion of 1.00g of propane?

3.7 – Limiting Reagents Limiting reactant or reagent Reactant that is completely consumed in a chemical reaction Theoretical Yield Amount of product you would get in a perfect experiment (determined mathematically) Actual yield What you actually get during the experiment (determined experimentally)

Example A strip of zinc metal massing 2.00g is placed in an aqueous solution containing 2.50g of silver nitrate, causing the following reaction to occur: Zn (s) + 2 AgNO 3 (aq)  2 Ag(s) + Zn(NO 3 ) 2 (aq) Which reactant is limiting? How many grams of Ag will form? How many grams of Zn(NO 3 ) 2 will form? How many grams of excess reactant will be left at the end of the reaction?

Example – Percent Yield Problem Imagine that you are working on ways to improve the process by which iron ore containing Fe 2 O 3 is converted into iron. In your tests you carry out the following reactions on a small scale: Fe 2 O 3 (s) + 3 CO (g)  2Fe(s) + 3CO 2 (g) If you start with 150g of Fe 2 O 3 as the limiting reagent, what is the theoretical yield of Fe? If the actual yield of Fe in your test was 87.9g, what is your percent yield?

2Al (s) + 3 Cl 2 (g)  2 AlCl 3 (s) A mixture of 1.50 mol Al and 3.00 mol Cl 2 are allowed to react. What is the limiting reagent? How many moles of AlCl 3 are formed? How many moles of excess reagent remain at the end of the reaction?