Chemistry Chapter on Density and Electrons

Significant Figures

• Multiplication/division: The final answer has the same number of significant figures as the measurement with the fewest significant figures.
• Addition/subtraction: The final answer has the same number of decimal places as the measurement with the fewest decimal places.

Unit Conversions

• Format: given unit x conversion factors = needed unit
• Example: Converting pounds to kilograms:
  • Given: 64 lbs
  • Needed: kg
  • Plan: lbs (US) --> kg (metric)
  • Conversion factors: 1 kg = 2.205 lb, or 2.205 lb/1 kg, and 1 kg/2.205 lb
  • Calculation: 64 lbs * (1 kg / 2.205 lb) = 29 kg

Density

• Density: the relationship of an object's mass to its volume (g/cm³, g/mL, or g/L)
• Every substance has a unique density.
• If an object is less dense than a liquid, it will float.
• Density of solids using volume displacement:
  • The volume of a solid can be determined by volume displacement; when a solid is submerged in water, it displaces a volume equal to its volume.

Problem-Solving Using Density

• Density can be used as a conversion factor.
• If the volume and density of a sample are known, the mass in grams can be calculated.

Classifying Substances

• Classify A-C as pure substances (element or compound) or mixtures (homogeneous or heterogeneous):
  • Copper wire: pure substance (element)
  • Chocolate chip ice cream: mixture (heterogeneous)
  • Nitrox: mixture (homogeneous)
• A salad dressing created with oil, vinegar, and chunks of cheese is a heterogeneous mixture.
• A mouthwash containing menthol, alcohol, hydrogen peroxide, and flavoring is a homogeneous mixture.

Physical and Chemical Changes

• Physical change: the physical properties of a substance change, but its identity stays the same.
  • Examples: boiling water, melting ice, freezing water, dissolving salt in water
• Chemical change: characteristics that indicate the ability of a substance to change into a new substance.
  • Examples: burning wood, cooking food, rusting of iron, digesting food, baking a cake.

States of Matter

• Solid: particles are tightly packed, resulting in a definite shape and volume.
• Liquid: flows freely, taking the shape of its container but retaining nearly constant volume.
• Gas: particles are widely spread, move randomly, and have neither definite shape nor volume.

Atomic Theory

• Dalton (1803): Proposed the atomic theory, which laid the foundation for modern chemistry
  • All matter is composed of tiny, indivisible particles called atoms
  • Atoms of the same element are identical in mass and properties, but atoms of different elements have different masses and properties
  • Atoms cannot be created or destroyed in a chemical reaction

Thomson (1897)

• Discovered the electron—the first subatomic particle.
• Used a cathode ray tube and the bending of particles towards a positive charge to prove the electron's negative charge.
• Proposed the plum pudding model of the atom.

Rutherford (1911)

• Discovered the nucleus—the atom's positively charged, dense center.
• Used the gold foil experiment to find the nucleus that deflected positively charges particles at large angles concluding that the atom has a small dense positively charged nucleus.

Bohr (1913)

• Contributed to the atomic theory by proposing the Bohr model.
• Electrons orbit the nucleus in distinct, stable energy levels.
• Electrons can jump between these energy levels when absorbing or emitting energy.

Isotopes

• Atoms of the same element have the same number of protons but different numbers of neutrons.

Atomic Emission Spectra

• A spectrum of the electromagnetic radiation emitted by a source.
• Electrons absorb energy, move to a higher energy level, and emit the energy in the form of a photon (producing color).
• Different colors correspond to different amounts of energy.

Electromagnetic Radiation

• Forms of energy, such as visible light, microwaves, X-rays, etc.
• Travel as waves (crest and trough) at the speed of light.
• Wavelength: distance between adjacent crests or troughs.
• Frequency: the number of times the crest passes a point in one second
• Frequency = energy
• Electromagnetic spectrum (arrangement of types of radiation)
• Period trends: ionization, electronegativity, atomic radius, metallic character, electronegativity

Periodic Trends

• Ionization energy: increases as you go up and to the right on the periodic table
• Electronegativity: increases as you go up and to the right on the periodic table
• Atomic radius: increases as you go down and to the left
• Metallic character: increases as you go down and to the left, decreases as you go up and to the right

Orbital Diagrams

• Represent the placement of electrons in orbitals
• Orbital configurations: notation that helps to show the placement

Pauli Exclusion Principle

• Each orbital can hold a maximum of two electrons with opposite spins.

Valence Electrons

• Determine how an atom will bond with other atoms
• Atoms try to achieve a full outer electron shell (usually 8 electrons, octet rule).
• Hydrogen and Helium need 2 electrons to fill their shell.

Electron Dot Diagrams (Lewis Dot Diagrams)

• Show valence electrons of an atom.
• In covalent bonds, pairs of electrons are shared
• In ionic bonds, one atom loses electrons (forming a cation), the other gains electrons (forming an anion).

Writing Formulas for Ionic Compounds

• Write the symbol for the cation (metal).
• Use appropriate oxidation state to show the charge.
• Write the symbol for the anion (non-metal).
• Look for charge and combine with metal.
• Balance charges.
• Balance ratio.

Naming Ionic Compounds

• Name the metal (cation) first.
• Name the nonmetal (anion) with the "-ide" suffix.
  • Use Roman numerals to denote charges for transition metals with multiple possible ion charges
• Combine the names
• Example: FeCl2 → Iron(II) chloride

Molecular Compounds

• Comprised of non-metals
• Use prefixes to indicate the number of each atom in the compound.
• Do not use mon- for first element
• Names end with -ide suffix.

VSEPR Theory (Valence Shell Electron Pair Repulsion)

• Predicts molecular shapes based on the repulsion of electron pairs around a central atom.
• Electron pairs repel each other, trying to get as far apart as possible.
• Electron pairs arranged to minimize repulsion.
  • Examples include linear (CO2), trigonal planar (BF3), tetrahedral (CH4), Trigonal Bipyramidal (PCl5), Octahedral (SF6).
• Lone pairs affect the arrangement and thus the shape.
• Water and Ammonia's molecular shapes are influenced by lone pairs.

Electronegativity and Polarity

• Electronegativity: the ability of an atom to attract electrons in a chemical bond.
• Polar covalent bond: electrons are shared unequally (one atom is more electronegative than the other). Uneven shared electrons create poles (partial positive and partial negative).
• Nonpolar covalent bond: electrons are shared equally.
• Polar molecules (e.g., water) have a positive and negative end, allowing them to dissolve other polar substances.