general chemistry 101 chem د. عبد العزيز بن علي الغامدي 1434-1435 هـ
TRANSCRIPT
References
• Chemistry the General Science 11E
T. L. Brown, H. E. LeMay, B. E. Bursten and C. J. Murphy.
العامة .الكيمياء . سليمان. د العويس، أحمد د
. الواصل، عبدالعزيز د الخويطر،السحيباني. عبدالعزيز .د
األول • الفصلي بعد: االختبار يحدد .لم
• : الثاني الفصلي بعد االختبار يحدد .لم
قريبا 134أ2المكتب • سيكون 146أ2و
[email protected]األيميل •
Matter
• Atoms are the building blocks of matter.• Each element is made of the same kind of atom.• A compound is made of two or more different kinds of
elements.
Units of Measurement
• SI Units
• Système International d’Unités• A different base unit is used for each quantity.
Metric SystemPrefixes convert the base units into units that are appropriate for the item being measured.
Volume
• The most commonly used metric units for volume are the liter (L) and the milliliter (mL).– A liter is a cube 1 dm long
on each side.– A milliliter is a cube 1 cm
long on each side.
Chemical Equations
Chemical equations are concise representations of chemical reactions.
CH4 (g) + 2 O2 (g) CO2 (g) + 2 H2O (g)
Anatomy of a Chemical Equation
CH4 (g) + 2 O2 (g) CO2 (g) + 2 H2O (g)
Reactants appear on the left side of the equation.
Anatomy of a Chemical Equation
CH4 (g) + 2 O2 (g) CO2 (g) + 2 H2O (g)
Products appear on the right side of the equation.
Anatomy of a Chemical Equation
CH4 (g) + 2 O2 (g) CO2 (g) + 2 H2O (g)
The states of the reactants and products are written in parentheses to the right of each compound.
Anatomy of a Chemical Equation
CH4 (g) + 2 O2 (g) CO2 (g) + 2 H2O (g)
Coefficients are inserted to balance the equation.
Subscripts and Coefficients Give Different Information
• Subscripts tell the number of atoms of each element in a molecule.
Examples of balancing chemical equation
HCl + Zn ZnCl2 + H2
C3H8 + O2 CO2 + H2O
Zn + HNO3 Zn(NO3)2 + H2 2
5 3 4
2
Combination Reactions
• In this type of reaction two or more substances react to form one product.
• Examples:– 2 Mg (s) + O2 (g) 2 MgO (s)
– N2 (g) + 3 H2 (g) 2 NH3 (g)
– C3H6 (g) + Br2 (l) C3H6Br2 (l)
Decomposition Reactions
• In a decomposition one substance breaks down into two or more substances.
• Examples:– CaCO3 (s) CaO (s) + CO2 (g)
– 2 KClO3 (s) 2 KCl (s) + O2 (g)
– 2 NaN3 (s) 2 Na (s) + 3 N2 (g)
Combustion Reactions
• These are generally rapid reactions that produce a flame.
• Most often involve hydrocarbons reacting with oxygen in the air.
• Examples:– CH4 (g) + 2 O2 (g) CO2 (g) + 2 H2O (g)
– C3H8 (g) + 5 O2 (g) 3 CO2 (g) + 4 H2O (g)
Formula Weights(FW)• A formula weight is the sum of the atomic
weights for the atoms in a chemical formula.
• So, the formula weight of calcium chloride, CaCl2, would be
Ca: 1(40.1 amu*) + Cl: 2(35.5 amu)
111.1 amu
• Formula weights are generally reported for ionic compounds.
*atomic mass unit
Molecular Weight (MW)• A molecular weight is the sum of the
atomic weights of the atoms in a molecule.
• For the molecule ethane, C2H6, the molecular weight would be
C: 2(12.0 amu)+ H: 6(1.0 amu)
30.0 amu
Percent Composition
One can find the percentage of the mass of a compound that comes from each of the elements in the compound by using this equation:
% element =(number of atoms)(atomic weight)
(FW of the compound)x 100
Percent Composition
So the percentage of carbon in ethane is…
%C =(2)(12.0 amu)
(30.0 amu)
24.0 amu
30.0 amu= x 100
= 80.0%
Avogadro’s Number
• 6.02 x 1023
• 1 mole of 12C has a mass of 12 g.• 1 mole of H2O has a mass of 18g.
Molar Mass
• By definition, a molar mass is the mass of 1 mol of a substance (i.e., g/mol).– The molar mass of an element is the mass
number for the element that we find on the periodic table.
– The formula weight (in amu’s) will be the same number as the molar mass (in g/mol).
Mole Relationships
• One mole of atoms, ions, or molecules contains Avogadro’s number of those particles.
• One mole of molecules or formula units contains Avogadro’s number times the number of atoms or ions of each element in the compound.
Examples
1- Calculate how many atoms there are in 0.200 moles of copper. The number of atoms in one mole of Cu is equal to the Avogadro number = 6.02 x 1023 . Number of atoms in 0.200 moles of Cu = (0.200 mol) x (6.02x1023 mol-1 ) = 1.20 x 1023 .
2- Calculate how molecules of H2O there are in 12.10 moles of water. Number of water molecules = (12.10 mol)x(6.02 x 1023) = 7.287 x 1024
3- Calculate the number of moles of glucose (C6H12O6) in 5.380 g of C6H12O6.
Moles of C6H12O6 = = 0.02989 mol.
4- Calculate the mass, in grams, of 0.433 mol of Ca(NO3)2.
Mass = o.433 mol x 164.1 g/mol = 71.1 g.
5.380 g
180.0 gmol-1
5- How many glucose molecules are in 5.23 g of C6H12O6? How many oxygen atoms are in this sample?
Molecules of C6H12O6 = x (6.02x1023)
= 1.75 x 1022 molecules
Atoms O = 1.75 x 1022 x 6 = 1.05 x 1023
5.23 g180.0 gmol-1
Calculating Empirical Formulas
One can calculate the empirical formula from the percent composition.
Calculating Empirical Formulas
The compound para-aminobenzoic acid is composed of carbon (61.31%), hydrogen (5.14%), nitrogen (10.21%), and oxygen (23.33%).
Find the empirical formula of PABA.
Calculating Empirical Formulas
Assuming 100.00 g of para-aminobenzoic acid,
C: 61.31 g x = 5.105 mol C
H: 5.14 g x = 5.09 mol H
N: 10.21 g x = 0.7288 mol N
O: 23.33 g x = 1.456 mol O
1 mol12.01 g
1 mol14.01 g
1 mol1.01 g
1 mol16.00 g
Calculating Empirical FormulasCalculate the mole ratio by dividing by the smallest number of moles:
C: = 7.005 7
H: = 6.984 7
N: = 1.000
O: = 2.001 2
5.105 mol0.7288 mol
5.09 mol0.7288 mol
0.7288 mol0.7288 mol
1.458 mol0.7288 mol
Combustion Analysis
• Compounds containing C, H and O are routinely analyzed through combustion in a chamber like this.– C is determined from the mass of CO2 produced.
– H is determined from the mass of H2O produced.– O is determined by difference after the C and H have
been determined.
Determining Empirical Formula by Combustion Analysis
*Combustion of 0.255 g of isopropyl alcohol produces 0.561 g of CO2 and 0.306 g of H2O. Determine the empirical formula of isopropyl alcohol.
Grams of C = x 1 x 12 gmol-1 = 0.153 g 0.561 g CO2 44 gmol-1
Grams of H = x 2 x 1 gmol-1 = 0.0343 g 0.306 g H2O 18 gmol-1
Grams of O = mass of sample – (mass of C + mass of H) = 0.255 g – ( 0.153 g + 0.0343 g) = 0.068 g
Moles of C = = 0.0128 mol 0.153 g C 12 gmol-1
Moles of H = = 0.0343 mol
Moles of O = = 0.0043 mol
0.0343 g H 1 gmol-1
0.068 g O 16 gmol-1
Calculate the mole ratio by dividing by the smallest number of moles:
C:H:O2.98:7.91:1.00
C3H8O
Stoichiometric Calculations
The coefficients in the balanced equation give the ratio of moles of reactants and products.
Stoichiometric Calculations
Starting with the mass of Substance A you can use the ratio of the coefficients of A and B to calculate the mass of Substance B formed (if it’s a product) or used (if it’s a reactant).
Stoichiometric CalculationsC6H12O6 + 6 O2 6 CO2 + 6 H2O
Starting with 1.00 g of C6H12O6… we calculate the moles of C6H12O6…use the coefficients to find the moles of H2O…and then turn the moles of water to grams.
Limiting Reactants
• The reactant that is completely consumed in a reaction is called the limiting reactant (limiting reagent). – In other words, it’s the reactant that run out first (in this
case, the H2).
Theoretical Yield
• The theoretical yield is the maximum amount of product that can be made.
• The amount of product actually obtained in a reaction is called the actual yield.
Theoretical Yield
• The actual yield is almost always less than the theoretical yield.
Why?• Part of the reactants may not react.• Side reaction.• Difficult recovery.
Percent Yield
• The percent yield of a reaction relates to the actual yield to the theoretical (calculated) yield.
Actual YieldTheoretical Yield
Percent Yield = x 100
Examples
Fe2O3(s) + 3 CO(g) 2 Fe(s) + 3 CO2(g)
If we start with 150 g of Fe2O3 as the limiting reactant, and found actual yield of Fe was 87.9 g, what is the percent yield?
The percent yield = x 100Actual Yield
Theoretical Yield
Theoretical yield = 105 g.
The percent yield = x 100 = 83.7 %
Fe2O3(s) + 3 CO(g) 2 Fe(s) + 3 CO2(g)
150 g Fe2O3
150 g 159 g mol-1
0.943 mol Fe2O32 mol Fe1 mol Fe2O3
X 1.887 mol Fe
1.887 mol x 55.85 g mol-1
105 g Fe
87.9 g105 g
Solutions
• Solutions are defined as homogeneous mixtures of two or more pure substances.
• The solvent is present in greatest abundance.
• All other substances are solutes.
Molarity
• Two solutions can contain the same compounds but be quite different because the proportions of those compounds are different.
• Molarity is one way to measure the concentration of a solution.
moles of solutevolume of solution in liters
Molarity (M) =
Mixing a Solution
• To create a solution of a known molarity, one weighs out a known mass (and, therefore, number of moles) of the solute.
• The solute is added to a volumetric flask, and solvent is added to the line on the neck of the flask.
Examples
Calculate the molarity of a solution made by dissolving 0.750 g of sodium sulfate (Na2SO4) in enough water to form 850 mL of solution.
Moles of Na2SO4 = = 0.0053 mol
Molarity = = 0.0062 M
0.750 g 142 g mol-1
0.0053 mol0.850 L
How many moles of KMnO4 are present in 250 mL of a 0.0475 M solution?
M =
Moles of KMnO4 = 0.0475 x 0.25 L
= 0.012 mol
molL
molL
How many milliliters of 11.6 M HCl solution are needed to obtain 0.250 mol of HCl?
moles of solutevolume of solution in liters
Molarity (M) =
volume in liters = 0.022 L = 22 mL.
11.6 M = 0.250 mol volume in liters
What are the molar concentrations of each of the ions present in a 0.025 M aqueous solution of Ca(NO3)2?
Ca2+ = 0.025 MNO3
- = 0.025 x 2 = 0.05 M
Dilution• One can also dilute a more concentrated solution
by– Using a pipet to deliver a volume of the solution to a
new volumetric flask, and– Adding solvent to the line on the neck of the new flask.
DilutionThe molarity of the new solution can be determined from the equation
Mc Vc = Md Vd
where Mc and Md are the molarity of the concentrated and dilute solutions, respectively, and Vc and Vd are the volumes of the two solutions.
Examples
How many milliliters of 3.0 M H2SO4 are needed to make 450 mL of 0.10 M H2SO4 ?
Mc Vc = Md VV
3.0 M x Vc = 0.10 M x 450 mL
Vc = 15 mL
Ways of Expressing Concentrations of Solutions
mass of A in solutiontotal mass of solution 100Mass % of A =
• Mass Percentage, ppm, and ppb Mass Percentage
Example: 36% HCl by mass contains 36 g of HCl for each 100 g of solution (64 g H2O)
Parts per Million (ppm)
ppm =
Parts per Billion (ppb)
mass of A in solutiontotal mass of solution
106
ppb = mass of A in solutiontotal mass of solution
109
Examples
Calculate the mass percentage of Na2SO4 in a solution containing 10.6 g Na2SO4 in 483 g water.
= 2.15 %
Mass % of Na2SO4 =10.6 g
(483 + 10.6) g 100
An ore contains 2.86 g of silver per ton of ore. What is the concentration of silver in ppm?
= 2.86 ppm
ppm = 2.86 g106 g
106
• Mole Fraction, Molarity, and Molality
Mole Fraction (X)
Molarity (M)
moles of Atotal moles in solutionXA =
moles of solutevolume of solution in liters
Molarity (M) =
Molality (m)
moles of solute Kilograms of solventm =
Since both moles and mass do not change with temperature, molality (unlike molarity) is not temperature-dependent.
Since volume is temperature-dependent, molarity can change with temperature.
ExamplesAn aqueous solution of hydrochloric acid contains 36 % HCl by mass. (a) Calculate the mole fraction of HCl in the solution. (b) Calculate the molality of HCl in the solution.
Moles HCl = = 0.99 mol
Moles H2O = = 3.6 mol
36 g 36.5 g mol-1
64 g 18 g mol-1