CHEM 011 Course Information

CHEM 011 Course Information

• The lecturer is Dr R Mhlaba.
• The classes are held in N-Block.
• Dr. Mhlaba's office is #1005.
• Class is held on Tuesdays at 11:10 (2 periods) and Fridays at 7:30 (2 periods).
• Practicals are conducted by Dr GR Monama, office 1020 N-Block.
• Labs require a lab coat and closed shoes.
• Lab coats and goggles can be purchased from Mr Seleka @ 068 021 7329 for R220, including free goggles.
• Test dates are Friday, April 3rd, and Monday, April 25th.
• Exams are in May/June.

Chemistry Definition

• Chemistry is the study of the composition and structure of materials, coupled with how those materials change.

International System of Units

• All measurements are made in the metric system.
• The metric system is a decimal system.
• Units are multiples of ten of larger or smaller units.

Fundamental Units of Measurement

• Length is measured in meters (m).
• Mass is measured in kilograms (Kg).
• Time is measured in seconds (s).
• Temperature is measured in Kelvin (K).
• Amount of substance is measured in moles (mol).

Derived Units

• Derived units result from the SI base units.
• There is no SI base unit for area.
• Area is calculated as Length x Width.
• Area is expressed as (meter) x (meter) = (meter)², or m x m = m².

Decimal Multipliers

• Prefixes are used to designate various multiples or submultiples in the metric system.
• Commonly used decimal multipliers serve as SI prefixes.
• Mega (M) has a multiplication factor of 10^6.
• Kilo (k) has a multiplication factor of 10^3.
• Deci (d) has a multiplication factor of 10^-1.
• Centi (c) has a multiplication factor of 10^-2.
• Milli (m) has a multiplication factor of 10^-3.
• Micro (µ) has a multiplication factor of 10^-6.
• Nano (n) has a multiplication factor of 10^-9.
• Pico (p) has a multiplication factor of 10^-12.
• Prefixes modify the size of the unit by the corresponding multiplier.
• The prefix "kilo" indicates a multiplying factor of 10^3 (1000).
• One kilometer equals 1 x 10^3 meters.
• A kilometer is a unit of length equal to 1000 meters.
• 1 dm = 1 / 10 m = 0.1 m

Measurements in the Laboratory

• Common measurements in the laboratory include length, volume, mass, and temperature.

Length Measurements

• SI base unit for length is meter (m).
• More convenient units include centimeters (cm) and millimeters (mm).
• 1 cm = 10^-2 m = 0.01 m
• 1 mm = 10^-3 m = 0.001 m
• 1 m = 100 cm = 1000 mm
• 1 cm = 10 mm

Volume Measurements

• Volume has a derived unit of (distance)³.
• Volume of a room calculation: Length x Width x Height = Volume.
• m x m x m = m³
• 1 L = 1 dm³
• The liter is too large for convenient lab measurements.
• Milliliters (ml) and cubic centimeters (cm³) are convenient units of volume for lab work.
• 1 L = 1000 cm³ = 1000 ml
• A cm³ can be abbreviated as cc.

Mass Measurements

• The kilogram (kg) is the base unit for mass.
• Grams (g) are a more convenient unit for laboratory measurement.
• 1g = 1 / 1000 kg = 0.001 kg
• 1 kg = 1000 g

Temperature Measurements

• The SI base unit for temperature is Kelvin (K).
• The three temperature scales for measurement are Kelvin, Fahrenheit, and Celsius.
• Water freezes at 0°C, 273 K, and 32°F.
• Water boils at 100°C, 373 K, and 212°F.
• The relationship between Kelvin (Tk) and Celsius (tc) is Tk = tc + 273.
• Celsius can be derived from Kelvin using: tc = Tk - 273.
• The relationship between Fahrenheit (TF) and Celsius (tc) is: TF = (9°F/5°C)tc + 32°F
• Celsius can be derived from Fahrenheit using: tc = (5°C/9°F)(tF - 32°F)
• Converting 25 °C to Kelvin: Tk = 25 + 273 = 298 K
• Converting 25 °C to Fahrenheit: TF = (9°F/5°C)(25) + 32 °F = 77 °F

Virus Length Exercise

• A virus that is 20 nm long.
• Objects smaller than about 2 x 10^-7 m cannot be seen under an optical microscope.

Significant Figures Definition

• Exact numbers have no uncertainty.
• Numbers derived from definitions; for example, 1 inch = 2.54 cm.
• Numbers are obtained from direct counting, like counting people in a room.
• Digits of a measured number that have uncertainty only in the last digit.
• The number of significant figures in a measured value is equal to the number of certain digits plus one uncertain digit.
• 11.2 ft has 3 significant figures.
• 11.13 ft has 4 significant figures.

Rounding

• Rounding is dropping insignificant digits.
• Rounding also involves adjusting the last digit reported in the final result of a calculation.

Rounding Procedure.

• Look at the leftmost digit to be dropped.
• IF the digit is 5 or greater:
  • Add 1 to the last digit to be retained
  • Drop all digits further to the right
• IF this digit is less than 5:
  • Drop all digits further to the right

Counting Significant Figures

• When zeros appear at the end of a number and to the right of the decimal point, they are counted as significant (e.g., 4.500 → 4 significant figures).
• Any zeros that precede the first nonzero digit are never counted as significant.
• 0.0023 has 2 significant figures.
• 45000 = 4.5 x 10^4 has 2 significant figures.
• 4.500 x 10^4 has 4 significant figures.
• Zeros at the end of a number are significant only if the decimal point is specified.
• 300. has 3 significant figures.
• 300.0 has 4 significant figures.
• Trailing zeros in numbers without a decimal point are not significant.
• 2050 has 3 significant figures.
• 300 has 1 significant figure.
• 300.0 has 4 significant figures.
• 100, 000, 000 has 1 significant figure.

Multiplication and Division with Sig Figs

• The answer should not have more significant figures than the least precise measurement used in calculation.
• (3.14 x 2.751)/ 0.64 = 13
• Calculator result is 13.497093
• Area of a rectangle with sides 12.34 cm and 1.23 cm is 15.2 cm².
• The calculator result is 15.1782 cm².

Addition and Subtraction with Sig Figs

• The answer should have the same number of decimal places as the quantity with the fewest decimal places
• 3.247 + 41.36 + 125.2 = 169.8
• Calculator result is 169.807
• 27.87 g - 21.2342 g = 6.64 g (2 decimal places)

Scientific Notation

• Scientific notation represents a number in form A × 10^n
• 1 ≤ A < 10, where n is an integer.
• Every digit included in A is significant.
• 0.000653 in scientific notation: 6.53 × 10^-4
• 350,000 in scientific notation: 3.5 × 10^5
• 0.02700 in scientific notation: 2.700 × 10^-2

Measurements And Significant Figures

• 310.0 kg has 4 significant figures.
• 0.224800 m has 6 significant figures.
• 0.05930 kg has 4 significant figures.
• 4.380 x 10^-8 m has 4 significant figures.
• 3.100 s has 4 significant figures.
• 91,000 has 2 significant figures.

Accuracy and Precision

• Accuracy is how close a measurement is to the actual, true value.
• Precision refers to how close multiple measurements are to each other.
• All measurements consist of a number, a unit, and a degree of uncertainty.
• Recurring a particular measurment usually leads to a slightly different result.
• This is because each measurement is prone to experimental error.

Unit Conversions (Factor Label Method)

• This is a method used to convert one system of units to another.
• (Given quantity) x (Conversion factor) = (Desired quantity)
• 72.0 inches is equal to 183 cm (1 inch = 2.54 cm)
• Conversion: 3.25 m = 3250 mm = 3.25 x 10^3 mm.
• Conversion: 29.21 ft = 8.903m
• Conversion details: Feet → Inches → centimetres → meters
• 29.21 ft x (12 inch / 1 ft) x (2.54 cm / 1inch) x (1 m / 100 cm) = 8.903m

Matter and Energy

• Matter: The atoms and molecules everything is made of.
• Mass: The quantity of matter an object has.
• Weight: The force with which the object is attracted by gravity.
• Physical states of matter: solid, liquid, and gas.
• Physical properties: are those that can be observed without changing the composition of a substance
• Chemical properties: are those that describe the chemical changes that a substance undergoes.
• Intensive properties: Those that are independent of the sample size.
• Examples of intensive properties: Color, electrical conductivity, etc.
• Extensive properties: Those that depend on the sample size.
• Examples of extensive properties: Mass and volume.

Chemical and Physical Changes

• Chemical Change = Chemical Reaction.
• A change in which one or more kinds of matter are transformed into a new kinds of matter.
• Examples of chemical changes: Rusting, Burning.
• When a material changes without changing its chemical identity then it is undergoing a physical change.
• Examples of physical changes: Physical state, Boiling point, Color
• Chemical changes are marked by the ability to react with other chemicals.
• Examples: Ability to react with oxygen, Ability to react with fluorine

Substance

• A kind of matter that cannot be separated into other kinds of matter by any physical process.
• Distillation and sublimation are two examples of physical processes.

Potassium Characteristics

• Potassium is a soft silvery-colored metal that melts at 64°C.
• It reacts vigorously with water, with oxygen, and with chlorine.
• Physical properties: Soft, Silvery , Melting point of 64 °C
• Chemical properties: React with water, React with oxygen, React with chlorine

Energy

• When chemical reactions occur they are accompanied by absorption and release of energy.
• Energy is the ability to do work.
• Two types of energy: Kinetic and Potential.
• Kinetic energy (KE): The energy an object has when it is moving.
• Potential energy (PE): Stored energy.
• Chemicals possess potential energy known as chemical energy.
• SI base unit for energy is joule (J = kg m² s^-2).
• calorie (cal) is another unit of energy
• 1 cal = 4.184 J or 1 kcal = 4.184 kJ

Density and Specific Gravity

• Density is defined as the ratio of an object’s mass to its volume.

Density Equation

• Density (D) = Mass (m) / Volume (V)
• Each pure substance has its own characteristics density.
• 47.3 cm³ has a mass of 37.32 g. Its density its is 0.789 g/cm³.
• The mass is 206 g given that a sample of vegetable oil has a density of 0.916 g/ml and the sample of volume is 225 ml.
• 130 ml is the volume of liquid needed given that 103 g of ethanol with a density of 0.789 g / ml, is needed for a chemical reaction.
• Specific gravity of a substance is defined as the ratio of the density of a substance to the density of water.

Specific Gravity Equation

• Specific gravity = Density of substance/ Density of water
• The density of pure water is 1.00 g/ml
• The specific gravity of methanol, with a density of 0.792 g/ml, is 0.792

Exercises

Perform the following calculations involving measurements and round the results to the correct number of significant figures.

• (i) 21.02333 g + 21.0 g
• (ii) 1.03 m x 2.074 m x 3.9 m
• (iii) 43.4 in x (1 ft / 12 in)

Solving Density of Blood

• A sample of blood completely fills an 8.20 cm³ vial.
• The empty vial has a mass of 10.30 g
• The vial has a mass of 18.91 g when filled with blood
• Calculating the density of blood requires the above parameters.
• Requires the application of mass and volume relations.