General Chemistry

Course Information

• Course: HNDT 1001 - General Chemistry for Nutrition
• Lecturer: Keisha Mascoll
• Semester: 1 2024/2025

Class Attendance

• All students must scan the QR code and complete required fields for class attendance.
• Download a QR code reader app (App Store or Google Play Store) if device is incompatible.

Course Meeting Times and Location

• Lecture: Thursday 1:00 PM - 3:00 PM - MSR3
• Lab: Thursday 3:00 PM - 5:00 PM - MLTC (lab)
• Tutorial: Friday 9:00 AM - 10:00 AM - MSR3

Course Rubric

• Mode of Delivery: Face-to-face
• Laboratory Session: Face-to-face
• Labs: 5, 25% weight, 5 weekly written reports on lab work assignments
• Quizzes: 2, 20% weight, multiple choice, short answer, & calculations. 1.5 hours
• Mid-Term Test: 1, 15% weight, multiple choice, short answer, & calculations, 1.5 hours
• Final Examination: 1, 40% weight, multiple choice, short answer, & calculations, 2 hours

Test Dates

• Quiz 1: Friday, October 4th, 2024, 1 hour, closed book
• Mid-term: Thursday, October 24th, 2024, 1 hour 30 minutes, closed book
• Quiz 2: Thursday, November 14th, 2024, 1 hour, closed book

Laboratory Schedule

• Week 4, Lab #1: Mole Ratio & Reaction Stoichiometry, Date: 26th September 2024
• Week 8, Lab #2: Titration of Acids & Bases, Date: 24th October 2024
• Week 10, Lab #3: Determination of Equilibrium Constant, Date: 7th November 2024
• Week 11, Lab #4: Calorimetry, Date: 14th November 2024
• Week 12, Lab #5: Worksheet, Date: 21st November 2024

Required/Essential Text

• Blackman, A., et al. (2019). Chemistry (4th ed.). New Jersey: John Wiley & Sons.
• ISBN-13: 978-0730363268, ISBN-10: 0730363287
• McMurry, J., et al. (2016). Fundamentals of General, Organic and Biological Chemistry (8th ed.). London: Pearson. ISBN-13: 978-0134015187, ISBN-10: 0134015185
• Petrucci, R., et al. (2016). General Chemistry: Principles and Modern Applications (11th ed.). London: Pearson. ISBN-13: 978-0132931281, ISBN-10: 0132931281

Learning Outcomes

• Describe the atom's structure (protons, neutrons, electrons), relative charges, and masses.
• Use chemical symbols to express isotopes' subatomic compositions, using atomic numbers and mass numbers.
• Describe the periodic table's organization by atomic number and chemical behavior (periods and groups).
• Identify locations of metals and nonmetals in the periodic table.
• Review electronic configurations.
• Differentiate between molecular and ionic substances.
• Explain ion formation (electron gain/loss) and predict common ion charges.

Topics

• Atomic Number & Mass Numbers
• Isotopes
• Periodicity
• Chemical Bonding
• Covalent Bonding
• Ionic Bonding
• Metallic Bonding
• Valence
• Octet Rule
• Ionic vs. Covalent Bonding

Atomic & Molecular Prospective

• Matter: atoms, ions, or molecules.
• Atom: smallest part of an element or compound that participates in chemical reactions.
• Molecule: groups of atoms held together with a specific connectivity and shape.
• Matter classified by: State (gas, liquid, solid), Composition (element, compound, mixture).

Elements, Compounds & Mixtures

• Pure Substance: fixed composition and distinct properties.
• Elements: identical atoms, e.g., oxygen (O2), gold (Au), silicon (Si), diamond (C).
• Compounds: more than one type of atom bonded, e.g., water (H2O), ethanol (C2H6O), sodium chloride (NaCl).
• Mixtures: variable composition, separable into components, properties depending on component amounts.

Molecular View of Elements & Compounds

• Classification of elements and compounds.
• Diagram showing atomic and molecular views of elements and compounds (e.g. Ne, O2, H2O, NaCl).

Homogeneous & Heterogeneous Mixtures

• Homogeneous Mixture: uniform throughout (e.g., air, coffee, vodka, blood, brass).
• Heterogeneous Mixture: non-uniform throughout (e.g., smog, ice cubes in drink, sand in water, concrete).

Structure of the Atom

• Protons (positive charge, mass 1), Neutrons (neutral charge, mass 1), Electrons (negative charge, negligible mass).
• Protons and neutrons located in nucleus, electrons orbit the nucleus.

Atomic & Mass Numbers

• Mass Number: sum of protons and neutrons.
• Atomic Number: number of protons (or electrons).

Isotopes

• Isotopes: atoms of the same element with different masses due to different neutron numbers.

Uses of Isotopes

• Stable isotopes used in nutritional studies.
• Practical applications: calcium and iron absorption, diet, physical activity, aging, medical therapy, nutrient metabolism, pregnancy energy cost, growth faltering, childhood and adult obesity.
• Isotopes utilized for metabolic studies, body composition assessment, and protein turnover.

Periodicity

• Elements arranged in periodic table by increasing atomic number, similar properties grouped together.
• Major classifications of elements: Main group (representative), transition metals, metalloids.

Major Classifications of Elements

• Main group elements: alkali metals, alkaline earth metals, halogens, noble gases.
• Transition metals.
• Metalloids.

Trends in the Periodic Table

• Trends of atomic radius, electron affinity, ionization energy, electronegativity.
• Trends across periods and groups.

Relative Abundance of Elements in the Human Body

• Human body comprised of ~ 99% of 6 elements (O, C, H, N, Ca, P) and a further 5 elements (S, K, Na, Cl, Mg), with ~0.15% of the rest being trace elements.

The Electronic Configuration of Atoms

• Aufbau Principle: electrons fill orbitals of lowest energy levels first.
• Pauli Exclusion Principle: an orbital can hold a maximum of 2 electrons with opposite spins.
• Hund's Rule: orbitals with the same energy are singly occupied before pairing occurs.

Shapes and Symmetry of Atomic Orbitals

• Atomic orbitals have specific shapes and symmetry.
• S orbital: spherical symmetry.
• P orbital: dumbbell shape, with two lobes.
• D orbital: more complex shapes.

Chemical Bonds

• Bond: force of attraction between oppositely charged particles.
• Covalent Bond: sharing of valence electrons between atoms.
• Ionic Bond: complete transfer of electrons from one atom to another.
• Metallic Bond: overlap of orbitals extending over three dimensions (3D) with delocalized electrons.

Covalent Bonding

• Sharing of valence electrons between atoms.
• Resulting molecules are product of covalent bonds, typically occurring between nonmetals.
• Different types of covalent bonds: single, double, triple bonds.

Properties of Covalent Compounds

• Low melting and boiling points.
• Poor conductors of electricity.
• Generally not soluble in water.

Ionic Bonding

• Complete transfer of electrons between a metal and a nonmetal.
• Formation of ions (cations (+) and anions (-) )
• Electrostatic force of attraction between oppositely charged ions.

Properties of Ionic Bonding

• High melting points.
• Poor electrical conductivity in solid state.
• Conduct electricity in molten or aqueous solution.

Metallic Bonding

• Overlapping of atomic orbitals in three dimensions.
• Delocalized electrons throughout the structure.
• Metals generally lose electrons to become positively charged cations.
• Tend to form ionic compounds with nonmetals.

Octet Rule

• Main group elements tend to form compounds by either gaining or losing electrons to achieve 8 valence electrons.

Exceptions to the Octet Rule

• Electron-deficient molecules: central atom has fewer electrons than normal.
• Hypervalent molecules: central atom has more electrons than normal.
• Odd-electron molecules: odd number of valence electrons, so electron is unpaired.

Ionic Bonding & Covalent Bonding

• Few compounds are 100% ionic or covalent.
• Ionic character depends on electronegativity difference.
• Covalent character in ionic compounds influenced by cation/anion size and charge.

Review Topics

• Atomic Structure & Composition, Periodic Table and its classifications.
• Characteristics of Metals, Non-metals, and Metalloids, Chemical Formulas, Chemical Nomenclature.
• Common Monoatomic Ions, Common Polyatomic Ions, Formula of Compounds.