Chemistry

SESSION 3 - UPCAT REVIEW 2018

CHEMISTRY OVERVIEW

• Definition: Chemistry is the science that deals with the composition, structure, and properties of substances, along with the transformations that they undergo.

BRANCHES OF CHEMISTRY

• Organic Chemistry: The study of compounds that contain both carbon and hydrogen.

• Inorganic Chemistry: The study of materials that do not contain hydrogen bonds.

• Analytical Chemistry: Involves quantitative and qualitative analysis of the chemical composition of materials.

• Physical Chemistry: The application of theories and concepts in physics (e.g., thermodynamics) to the study of chemical systems.

MATTER

• Definition: Matter is a material substance that occupies space, has mass, and is composed predominantly of atoms (protons, neutrons, and electrons) and is interconvertible with energy.

STATES OF MATTER

• Solid: Holds shape and has a fixed volume.

• Liquid: Takes the shape of the container, has a free surface, and has a fixed volume.

• Gas: Fills the shape of the container and has no fixed volume.

PHASE CHANGES

• Sublimation: Transition from solid to gas.

• Deposition: Transition from gas to solid.

• Melting: Transition from solid to liquid.

• Evaporation: Transition from liquid to gas.

• Condensation: Transition from gas to liquid.

• Freezing: Transition from liquid to solid.

CHANGES IN MATTER

• Physical Change: Some physical properties of a substance are modified without changing its chemical composition.

• Chemical Change: Change in the composition of matter that produces a new substance with new properties.

CLASSIFICATION OF MATTER

• Pure Substances: Substances that contain only one kind of molecule, cannot be separated by physical means.

  • Elements: Cannot be decomposed into simpler substances.

  • Compounds: Composed of two or more elements (e.g., Water (H2O), Sodium Chloride (NaCl)).

• Mixtures: Can be separated into two or more substances by physical means.

  • Homogeneous Mixtures: Substances are evenly distributed (e.g., Blood, Air, Juice).

  • Heterogeneous Mixtures: Non-uniform composition (e.g., Oil and Water, Rocks, Pizza).

  • Suspensions: Solids immersed in liquid (e.g., Juice, Sand and Water).

  • Colloids: Particles are evenly distributed (e.g., Milk, Fog). They do not settle over time; exhibit Tyndall effect (light scattering).

  • Immiscible Liquids: Liquids that cannot mix (e.g., Oil and Water).

CHARACTERISTICS OF MATTER

• Physical Properties: Innate characteristics of matter (e.g., color, melting point, conductivity).

• Chemical Properties: Characteristics observed when the composition changes (e.g., flammability, toxicity).

• Intensive Properties: Independent of the amount of substance (e.g., color, density).

• Extensive Properties: Dependent on the amount of substance (e.g., mass, length).

MEASUREMENTS IN MATTER

• S.I. Units: Le Systeme International d’ Unites or the International System of Units.

  • Mass: kilogram (kg)

  • Length: meter (m)

  • Time: seconds (s)

  • Amount of Substance: moles (mol)

  • Electric Current: Ampere (A)

  • Luminous Intensity: candela (cd).

DIMENSIONAL ANALYSIS

• Definition: The conversion of one unit to another.

• Example: Convert density from $13.6 ext{kg/cm}^3$ to $ext{kg/m}^3$. Calculation steps:

  • 
    ho ext{ in } ext{kg/m}^3 = 13.6 ext{g} imes rac{1 ext{cm}^3}{1 ext{cm}^3} imes igg( rac{100 ext{cm}}{1 ext{m}}igg)^3 imes rac{1 ext{kg}}{1000 ext{g}}

  • Result: $1.36 imes 10^4 ext{ kg/m}^3$.

COMPONENTS OF MATTER

• Discovery of Electrons: Early experiments with Cathode Ray Tubes by J.J. Thomson that revealed electrons.

• Oil Drop Experiment: Conducted by Robert Milikan to determine the charge and mass of an electron.

• Structure of Matter:

  • Rutherford's Gold Foil Experiment: Introduced the concept of a nucleus and protons.

  • Chadwick: Discovered neutrons (neutrally charged particles).

DALTON’S ATOMIC THEORY

1. All matter consists of tiny, indivisible particles called "atoms".

2. Atoms of an element cannot be created, destroyed, or transformed into another element.

3. Atoms of the same element are identical in mass and other properties and are different from atoms of any other element.

4. Compounds result from the chemical combination of specific ratios of atoms of different elements.

LAWS OF CHEMICAL COMBINATION

• Law of Conservation of Mass: Mass is neither created nor destroyed in chemical reactions.

• Law of Definite Proportions: In a given compound, the constituent elements are always combined in the same proportions by mass.

• Law of Multiple Proportions: The masses of one element that combine with a fixed mass of the second element can be expressed as ratios of small whole numbers.

ATOMS AND ELEMENTS

• Mass Number (A): $A = ext{No. of PROTONS} + ext{No. of NEUTRONS}$.

• Atomic Number (Z): $Z = ext{No. of PROTONS}$.

• Relationship: No. of PROTONS = No. of ELECTRONS (neutral atoms).

ISOTOPES

• Elements with the same atomic number but different atomic weights.

ATOMIC WEIGHT & MOLECULAR WEIGHT

• Atomic Weight: Weighted average of the atomic masses of the isotopes of an element.

• Molecular Weight: Sum of the atomic weights of atoms in a molecule.

  • Example: For glucose $C<em>6H</em>{12}O_6$:

  • 6 C = $6 imes 12.011 ext{ amu} = 72.066 ext{ amu}$.

  • 12 H = $12 imes 1.0079 ext{ amu} = 12.095 ext{ amu}$.

  • 6 O = $6 imes 15.9994 ext{ amu} = 95.9964 ext{ amu}$.

  • Total: 180.097 amu.

• Molar Mass: A mole of any substance has a mass in grams equal to its atomic weight.

• Avogadro’s Number: Constant of $6.022 imes 10^{23}$ particles per mole.

ELECTRON CONFIGURATION

• Definition: Distribution of electrons into energy levels and sublevels that determines the behavior of elements.

• Aufbau Principle: Electrons are added successively to the lowest energy orbitals.

• For ground state atoms:

  • Electrons occupy energy shells, subshells, and orbitals that minimize energy.

  • Sequence: Start from 1s and move upwards.

QUANTUM NUMBERS

• Principal Quantum Number (n): Values from 1, 2, 3…

• Angular Momentum Number (l): Values from 0 to (n-1), indicating subshells (s, p, d, f).

• Magnetic Quantum Number (m₁): Values that indicate the orientation of the orbital.

• Spin Quantum Number (m_s): Indicates the spin direction of electrons, either +1/2 or -1/2.

• Hund’s Rule: Every orbital in a subshell is singly occupied before any orbital is doubly occupied, and all electrons in singly occupied orbitals have the same spin.

• Pauli Exclusion Principle: No two electrons in the same atom can have the same four quantum numbers.

PERIODIC TRENDS

• Trends:

  • Atomic Size: Generally increases down a group and decreases across a period.

  • Ionization Energy: The energy required to remove an electron; increases across a period and decreases down a group.

  • Electron Affinity: Energy change when adding an electron; generally becomes more negative across a period.

  • Electronegativity: Ability of an atom to attract electrons in a bond; increases across a period and decreases down a group.

  • Metallic Behavior: Increases down a group and decreases across a period.

TYPES OF BONDS

1. Ionic Bond: Electrons are transferred between atoms, forming cations and anions.

2. Covalent Bond: Electrons are shared between atoms to form molecules.

3. Metallic Bond: A network of positive ions is immersed in a sea of delocalized electrons.

CHEMICAL REACTIONS

• Described by chemical equations with reactants on the left and products on the right.

• Types of Reactions:

  1. Combustion Reaction: Oxygen combines with another compound, forming water and carbon dioxide.

  2. Synthesis Reaction: Two or more reactants combine to form a more complex product.

  3. Decomposition Reaction: A complex molecule breaks down into simpler substances.

  4. Single Displacement Reaction: An element trades places with another element in a compound.

  5. Double Displacement Reaction: The anions and cations of two different molecules switch places.

  6. Acid-Base Reaction: The H⁺ ion from the acid reacts with the OH^- ion from the base, yielding salt and water.

FACTORS INFLUENCING CHEMICAL REACTIONS

• Nature of Reactants: More reactive reactants lead to faster reactions.

• Concentration of Reactants: Higher concentrations yield faster reactions.

• Temperature: Increased temperature accelerates reaction rates.

BALANCING CHEMICAL EQUATIONS

• Adjust coefficients to ensure the number of atoms of each element is equal on both sides of the equation.

  • Example: Balance the combustion of propane $C<em>3H</em>8 + O<em>2 ightarrow CO</em>2 + H_2O$.

• Two balanced examples:

  1. $Fe<em>2O</em>3 + C <br>ightarrow Fe + CO<em>2$ results in $2Fe</em>2O<em>3 + 3C ightarrow 4Fe + 3CO</em>2$.

  2. $HCN + NO<em>2 ightarrow C</em>2N<em>2 + NO + H</em>2O$ results in $2HCN + NO<em>2 ightarrow C</em>2N<em>2 + NO + H</em>2O$.

STOICHIOMETRY

• Using a balanced chemical equation to calculate the quantities of reactants and products.

• Ratios derived from stoichiometric coefficients indicate the relative proportions of substances in reactions.

• Example Calculation: To find grams of NaOH needed to react with H2SO4: $2NaOH(aq) + H<em>2SO</em>4(aq) <br>ightarrow 2H<em>2O + Na</em>2SO_4(aq)$ interprets to 2.53g of NaOH.

SOLUTIONS

• Solute: Substance that dissolves in a solvent (e.g., lower amount changing state).

• Solvent: Dissolving medium.

• Solution: Homogeneous mixture of solute(s) in a solvent.

KINETIC-MOLECULAR THEORY

1. Gases consist of many small particles moving in straight-line motion.

2. Gas molecules occupy no volume (considered point masses).

3. Collisions between gas molecules are perfectly elastic.

4. There are no attractive or repulsive forces between gas molecules except during collisions.

5. The average kinetic energy of a molecule is given by $rac{3}{2}kT$, where T is the absolute temperature and k is the Boltzmann constant.

GAS LAWS

• Pressure Units: Atm (atmospheres), mmHg, Torr, Pa (Pascals).

• Conversions:

  • Temperature: $K = °C + 273$.

  • Volume: $1 ext{ cm}^3 = 1 ext{ mL}$, $1 ext{ dm}^3 = 1 ext{ L} = 1000 ext{ mL}$.

• Standard Conditions:

  • $0.00 °C = 273 K$;

  • $1.00 atm = 760 mmHg$;

  • $1.00 atm = 101.325 kPa$.

BOYLE’S LAW

• Pressure and volume inversely related when temperature is constant:

  • $P<em>1V</em>1 = P<em>2V</em>2$.

  • Example calculation uses pressure conversion if units need to match.

CHARLES’ LAW

• Volume and temperature directly related when pressure is constant:

  • $rac{V<em>1}{T</em>1} = rac{V<em>2}{T</em>2}$.

  • Temperature must be converted to Kelvin for calculations.

COMBINED GAS LAW

• Combines Boyle's, Charles', and Gay-Lussac's Laws:

  • $rac{P<em>1V</em>1}{T<em>1} = rac{P</em>2V<em>2}{T</em>2}$.

AVOGADRO’S LAW

• Volume is proportional to the number of particles at constant temperature and pressure:

  • $rac{V<em>1}{N</em>1} = rac{V<em>2}{N</em>2}$.

IDEAL GAS LAW

• Combines gas laws into one equation:

  • $PV = nRT$; where R = 0.08206 atm·L/mol·K.

• Example calculation involves pressure conversion and application of the Ideal Gas Law.

DALTON’S LAW OF PARTIAL PRESSURES

• Each gas in a mixture exerts pressure as if other gases are absent:

  • $P<em>{ ext{total}} = P</em>1 + P<em>2 + P</em>3 + … + P_n$.

REFERENCES

• Petrucci, R., et al. (2011). General Chemistry: Principles and Modern Applications (10th ed.). Toronto, ON: Pearson Canada, Inc.

• Silberberg, M. (2013). Principles of General Chemistry (3rd ed.). New York, NY: McGraw-Hill Companies, Inc.

• Google Images.