Chem test

SNC2DE Unit: 3.0 Chemistry Lesson:

3.01 Topics: Introduction to Chemistry

A) Review of Safety, Lab, etc.

• Discussion points include:

  • Review of Hazardous Household Products (HHPs)

  • Overview of Workplace Hazardous Materials Information System (WHMIS) symbols and their meanings

  • Overview of lab equipment handout and brief lab tour

B) Classification of Matter

• Definition of Matter:

  • Anything that occupies space and has mass.

• Types of Matter:

  • Pure Substances

  • Consists of matter with uniform composition. Components cannot be separated by physical means.

  • Changes of state occur at constant temperature.

  • All particles have identical properties.

  • Elements

  • Cannot be broken down further, comprised of only one type of atom or molecule; approximately 114 elements known so far.

  • Atoms

    • The smallest unit of matter.

  • Compounds

  • Contain 2 or more elements in a fixed proportion. Breaking it down results in substances with different properties (e.g., $ext{NaCl} <br>ightarrow ext{Na} + ext{Cl}$).

  • Molecule

    • The smallest entity maintaining the same proportions of constituent atoms as the compound.

  • Mixtures

  • A mechanical combination of 2 or more pure substances. Each retains its properties and can be separated easily.

  • Heterogeneous Mixture

    • Different substances appear as different phases (e.g., mud, oil, and water).

  • Homogeneous Mixture (Solution)

    • Only one phase is visible (e.g., salt and water).

• States of Matter:

  • Matter can exist in three states: Solid (s), Liquid (l), or Gas (g).

• Changes of State:

  • Includes melting/freezing (fusion), boiling/condensing (vaporization), and sublimation.

  • Aqueous Solutions (aq): Solutions where water is the solvent (examples include acids and bases).

C) Physical and Chemical Properties and Changes

• Physical Properties:

  • Characteristics used to describe a substance:

  • Colour, lustre, ductility (ability to be drawn into a wire), malleability (ability to be hammered), texture, smell, solubility, taste, melting point, conductivity (thermal or electrical), density, crystalline/amorphous state, hardness, volatility, viscosity, and transparency.

• Physical Changes:

  • Changes that do not produce new substances, include:

  • Changes of state, physical separation, generally reversible, and involve no significant heat exchange.

• Chemical Properties:

  • Behavior when substances react to form new ones:

  • Flammability, bleaching ability, corrosion, reactions with acids or bases.

• Chemical Changes:

  • Produce one or more new substances and are usually irreversible.

• Indicators of Change Type:

  • Chemical:

  • Colour change, heat/light produced/absorbed, precipitate formed, gas formed, mass change.

  • Physical:

  • Change of state, dissolving, and mechanical actions such as cutting.

• Gas Tests:

  • Glowing Splint: Reignites in Oxygen.

  • Burning Splint: "Pop" sound indicates Hydrogen presence.

  • Lime Water Test: Turns cloudy if Carbon Dioxide is present.

  • Cobalt Chloride Paper: Turns from blue to white/pink in the presence of water vapour.

Notes:Classification of Matter

pure substance Same composition throughout; can't be separated physically; changes of state at constant temp

element Only ONE type of atom. Can't be broken down further (e.g. O, Na, Fe)

compound 2+ elements in a FIXED ratio. Breaking it down changes properties (e.g. NaCl → Na + Cl)

mixture 2+ pure substances physically combined. CAN be separated. Variable proportions.

heterogeneous Different phases visible (mud, oil & water)

homogeneous Only one phase visible = solution (salt water)

Physical Properties & Changes

Properties:colour, lustre, ductility, malleability, texture, smell, solubility, melting point, density, state, conductivity, hardness, viscosity

Changes: change of state, dissolving, cutting — NO new substance. Often reversible.

Chemical Properties & Changes

Properties: flammability, bleaching ability, corrosion, reaction with acids/bases

Changes: new substance(s) formed. Often irreversible.

Signs of Chemical Change (clues, not proof)

chemical clues Colour change · heat or light produced/absorbed · precipitate formed · gas produced · mass change*

*caution Boiling water produces gas but is PHYSICAL. These are clues, not definitive proof.

physical clues Change of state · dissolving · mechanical actions (cutting)

• Homework: Flow Chart for Classification of Matter, homework sheet C1

3.02 The Atom, Bohr and Lewis Diagrams

A) Introduction

• Basic structure of the atom: protons, neutrons, and electrons.

B) Atomic Structure

• Electrons have a negative charge and orbit the nucleus (which contains protons and neutrons).

• Protons (positive charge) and electrons (negative charge) have equal and opposite charges known as elementary charges.

• Neutrons possess no charge.

• Size Comparison: Protons/neutrons are about 2000 times larger than electrons.

• Atoms are mostly empty space.

  • Example: If a hydrogen nucleus were 1mm, the electron would orbit at a 100m diameter.

• Electrons reside in specific orbital shells (each shell can hold a maximum number of electrons: 2, 8, 8).

  • Electrons fill the innermost shells before working outward.

  • The final unfilled shell is called the valence shell (contains valence electrons).

• Formula for maximum electrons in each shell: $2n^2$ where n is the energy level number.

• Atomic Notation:

  • Example: $ext{A}_{ ext{Z}}^{ ext{B}}$ where A is the atomic number (number of protons), B is the atomic mass number (number of protons + neutrons).

  • Atomic mass unit (amu) is defined as follows:

  • For $^{12}C, 12$ amu, where 1 proton = 1 neutron = 1 amu.

C) Bohr and Lewis Diagrams

• Bohr Diagrams:

  • Displays important atomic data but can be cumbersome.

• Example for Magnesium: $ext{Mg}_{12}^{24}$

  • Simplified structure:

  • (12$p^+$, 12$n^0$) 2$e^-$, 8$e^-$, 2$e^-$ .

• Lewis Diagrams:

  • Displays only the valence electrons for a more concise representation.

  • Example:

  • Mg: ••

  • Al: •• ••

  • O: • •

  • Li: ∙ •• ••

  • Cl: •

• Emphasis on the importance of Lewis diagrams for visualizing molecules.

D) Isotopes

• Definition:

  • Isotopes are variations of elements with the same number of protons but different numbers of neutrons.

• Isotopes behave chemically alike yet can exhibit physical differences.

• Example: Carbon has isotopes $^{12}C$, $^{13}C$, and $^{14}C$ with relative abundances of 99%, 1%, and negligible, respectively.

  • $^{12}C$ and $^{13}C$ are stable, while $^{14}C$ is unstable and utilized in radiometric carbon dating.

E) Ions

• Neutral atoms contain equal numbers of protons and electrons.

• Elements form ions by gaining or losing electrons to achieve a full outer shell (octet rule).

• Ions:

  • Cations: Positively charged (loss of electron).

  • Anions: Negatively charged (gain of electron).

  • Nomenclature: Suffix ‘ide’ for anions (e.g., Fluorine to Fluoride).

• General trend:

  • Metals lose electrons to form cations.

  • Non-metals gain electrons to form anions.

• The process of gaining electrons releases energy (termed electron affinity energy).

3.03 Topics: The Periodic Table

A) Introduction

• Previous discussions focused on atomic structure; the periodic table offers systematic grouping based on elemental properties.

B) The Periodic Table

• Historical attempts to categorize elements date back to the beginnings of alchemy.

• Dmitri Mendeleev published the first periodic table in 1869, predicting undiscovered elements based on patterns observed in the elements studied.

• Elements are arranged in a modern periodic table into groups and periods for pattern recognition:

  • Groups (columns):

  • 18 groups, can be represented as 8 groups (ignoring transition metals) where groups III to VIII are equivalent to groups 13 to 18.

  • Elements in the same group exhibit similar valence electron configurations.

  • Periods (rows):

  • 7 periods, with each period representing a new shell for electrons.

  • Each element in the same period has the same number of shells.

• Significant Groups:

  • Group I - Alkali Metals: Shiny metals, highly reactive; compounds formed are basic oxides and mostly soluble in water.

  • Group II - Alkaline Earth Metals: Shiny metals; form mostly insoluble compounds.

  • Group VII - Halogens: Nonmetals, highly reactive, poisonous.

  • Group VIII - Noble Gases: Highly stable nonmetals with limited reactivity.

• Other Groups:

  • Transition metals: Dense, good conductors, malleable, generally solid at room temperate (except Mercury).

  • Metalloids: Elements that may behave as metals and nonmetals.

  • Nonmetals: Diverse states at room temperature, with common occurrences in the earth's crust and biology.

C) Atomic Radius

• The size of an atom is influenced by its electron shells and pairs.

• Atomic radius trends:

  • Decreases left to right across periods (greater attraction as protons/electrons increase).

  • Increases down a group (more orbitals increase size).

• Key Relationships:

  • Atomic radius decreases across a period.

  • Atomic radius increases down a group.

• Extensions:

  • 118 elements discovered; f-block elements exist outside of the main table.

• Periodic Table Blocks:

  • s-block: Groups I and II, consistent trends in properties based on electron configurations.

  • p-block: Last six groups excluding helium, containing all nonmetals and semimetals.

  • d-block: Groups 3-12, involving filling of d-orbitals.

  • f-block: Consists of lanthanides and actinides, involving f-orbitals.

D) Quantum Orbitals

• Each orbital has quantum numbers corresponding to electron energy, angular momentum, and orientation.

• Orbitals are characterized as follows:

  • s orbitals: ℓ = 0, spherical.

  • p orbitals: ℓ = 1, dumbbell shaped.

  • d orbitals: ℓ = 2, more complex shapes.

  • f orbitals: ℓ = 3, even more complex.

E) Summary of Orbital Filling Rules

• Aufbau Principle: Electrons fill orbitals from lowest to highest energy.

• Hund's Rule: Each orbital in a subshell is singly occupied before any is doubly occupied.

• Pauli Exclusion Principle: No two electrons can have the same set of all four quantum numbers in an orbital.

3.04 Topics: Trends in the Periodic Table

A) Introduction

• Previous lessons discussed elemental arrangement in periodic table; today's focus is on elemental behaviors and relations.

B) Reactivity

• Elements in the same group behave similarly due to shared valence electrons.

  • Metal reactivity is higher compared to non-metals.

  • Metal electrons are often given up easily compared to nonmetal electrons, which are attracted.

• Reactivity trends:

  • Increases down groups for metals.

  • Increases up groups for non-metals.

C) Ionization Energy (I.E.)

• Energy required to remove an electron from an atom.

• I.E. is lower for more reactive elements.

• Trends in I.E.:

  • Increases across periods (more tightly held electrons).

  • Decreases down groups (higher energy levels increase distance).

• Example Outcomes:

  • Compare I.E.s for Boron (800 kJ/mol) and Magnesium (738 kJ/mol) showing the increasing energy required for successive removals.

D) Electron Affinity

• A measure of an element's attraction for electrons.

• Trends in electron affinity:

  • Increase across a period (smaller atomic radius, stronger pull from nucleus).

  • Decrease down a group (shielding effect).

E) Electronegativity

• Attraction of an atom for electrons in a bond, evaluated using the (Linus) Pauling scale.

• Trends:

  • Increases across a period.

  • Decreases down a group.

F) Effective Nuclear Charge

• Net positive charge felt by valence electrons; calculated as protons - shielding inner electrons.

• Z increases across a period but decreases down a group due to increased distance of valence electrons.

G) Metallic Character

• Characteristic properties associated with metals (ductility, malleability, conductivity).

• Trends:

  • Decreasing metallic character across a period.

  • Increasing metallic character down a group.

3.05 Topics: Ionic Bonding and Ionic Compounds

A) Ionic Compounds

• Positive cations and negative anions form stable ionic compounds through electron transfer.

• Example: Sodium chloride $ext{NaCl}$ and its formation:

  • The Lewis structures depict the formation of ions aligning through charge interactions.

• Properties of Ionic Compounds:

  • High melting and boiling points, form crystals, dissolve in water, conduct electricity when dissolved.

  • Not soluble in non-polar solvents.

• Electronegativity Considerations:

  • Ionic bonds form when the electronegativity difference is >1.8.

  • Solid ionic compounds conduct electricity when dissolved in water but not in solid-state.

B) Lattice Structure of Ionic Compounds

• Coordination numbers indicate the number of cations or anions neighboring a particle.

• Common Lattice Structures:

  • Structures vary with ionic radii ratios affecting stability and bond strength.

C) Examples of Ionic Compounds:

• Formation of lattice structures in ionic compounds is stable due to opposite charges attracting each other, releasing lattice energy upon formation.

3.06 Topics: Spectra and Naming Conventions

A) Spectra

• Discussion focused on electron excitement to higher orbits leading to light emissions upon return to lower states.

B) Naming Formula for Binary Ionic Compounds

• Cross Over Method:

  • Charge crisscross method to establish binary compound formulas.

• Naming Conventions:

  • Metal first, nonmetal adopts “-ide” suffix for naming.

• Example: $ext{MgCl} <br>ightarrow ext{Magnesium Chloride}$.

C) Multiple Ionic Charges

• Transition metals can exhibit multiple oxidation states, thereby needing careful naming (e.g., Iron (II) vs Iron (III)).

3.07 Topics: Polyatomic Ions

A) Polyatomic Ions

• Groups of bonded atoms with a net charge.

• Commonly negatively charged, except ammonium (NH4^+).

• Naming follows from the charge (-ate suffix for most).

B) Naming Considerations

• Varying oxygen in polyatomic ions results in using prefixes (hypo-, per-) and suffix adjustments for clarity (e.g., Chlorate, Hypochlorite).

SNC2DE Unit: 3.0 Chemistry Lesson: 3.08 Topics: Naming of Binary Acids

A) Binary Acids

• Definition: Acids that include Hydrogen and another non-metal element.

• Naming format: Hydro + base name of the second element + ic + acid (e.g., HCl = Hydrochloric Acid).

B) Oxy-Acids

• Derived from polyatomic ions, naming similar to binary acids but with variations based on oxygen presence.

• Naming similar to polyatomic ion modifiers; H2SO4 = Sulfuric Acid from sulfate.

SNC2DE Unit: 3.0 Chemistry Lesson: 3.09 Topics: Covalent Bonding

A) Covalent Bonding

• Bonds formed between non-metals through shared pairs of electrons.

• Covalent bonds do not conduct electricity in solutions.

• Sharing arrangements are displayed using Lewis diagrams (e.g., Methane, Ammonia).

B) Properties of Covalent Compounds

• Non-ionic, varied molecular compounds in daily life, often hydrocarbons.

• Examples: Appropriately named based on prefixes or common names.

C) Electronegativity

• Applied in predicting bond types: Ionic (>1.7), Polar Covalent (0.7 - 1.7), Purely Covalent (<0.7) based on differences in electronegativity.

SNC2DE Unit: 3.0 Chemistry Lesson: 3.10 Topics: Molecular Compounds

A) Naming Covalent Compounds

• Use of prefixes to denote numbers of each element present.

• Example: $ext{PCl}_3 = ext{Phosphorus Trichloride}$.

B) Molecular Geometry

• Basic molecular geometry considerations, including VSEPR theory and bond angles.

• Visual representation techniques (Ball and Stick, Electron Density Plot) to indicate molecular structures.

SNC2DE Unit: 3.0 Chemistry Lesson: 3.11 Topics: Chemical Equations & Conservation of Mass

A) Chemical Reactions

• Focus on reactants to products and methods to describe these transitions using word and chemical equations.

B) Conservation of Mass

• Principle established by Antoine Lavoisier: Mass is preserved in chemical reactions, necessitating balanced equations for accurate representation.

  • Emphasizes distinction in coefficients vs subscripts when counting atoms in various formulas, ensuring accurate mass representation throughout reactions.

Notes:

• Homework: Begin C08.