Definitions & Examples

Definitions & Examples

• A pure substance has a uniform and definite composition, consisting of only one type of particle; water (H₂O) and oxygen (O₂) are examples.
• An element is a pure substance that cannot be broken down into simpler substances through chemical means; hydrogen (H) and iron (Fe) are examples.
• A compound is a pure substance made of two or more elements chemically combined in a fixed ratio; carbon dioxide (CO₂) and water (H₂O) are examples.
• An atom is the smallest unit of an element that retains its chemical properties; a single oxygen atom (O) is an example.
• Density is the amount of mass per unit volume of a substance, calculated as Density=Mass/Volume; water has a density of 1 g/cm³.
• A heterogeneous mixture has visible, non-uniformly distributed components; raisin bran cereal, salad, and sand with iron filings are examples.
• A homogeneous mixture (or solution) looks uniform throughout with evenly distributed particles; saltwater, air, and brass are examples.
• Isotopes are atoms of the same element with differing numbers of neutrons but the same number of protons; carbon-12 and carbon-14 are isotopes of carbon (C).
• Boiling point is the temperature at which a substance changes from liquid to gas; water boils at 100°C at standard atmospheric pressure.
• Freezing/melting point is the temperature at which a substance changes between solid and liquid states; ice melts at 0°C.
• An alloy is a mixture of two or more metals or a metal and another element; steel (iron and carbon) and brass (copper and zinc) are examples.
• Viscosity measures a fluid's resistance to flow; honey has high viscosity, while water has low viscosity.
• A combustible substance can catch fire and burn in the presence of oxygen; wood and gasoline are examples.
• Malleability is the ability of a substance to be hammered or rolled into thin sheets; gold is highly malleable.
• Ductility is the ability of a material to be stretched into a wire; copper is highly ductile.
• Lustre is how a material reflects light, either metallic or non-metallic; silver has a shiny metallic lustre.
• Optical clarity measures a material's ability to transmit light; glass is transparent, while frosted glass is translucent.
• Solubility is the ability of a substance to dissolve in a solvent; sugar dissolves in water, but sand does not.
• Brittleness is the tendency of a material to break or shatter under stress; glass is brittle, while rubber is not.

The Five Points of the Particle Theory

• All matter is made up of tiny particles (atoms or molecules).
• Particles are always in motion, vibrating in place in solids, moving more freely in liquids, and rapidly in all directions in gases.
• Particles attract each other, with the strength of attraction strongest in solids, weaker in liquids, and weakest in gases.
• Particles have spaces between them, smallest in solids, larger in liquids, and largest in gases.
• Adding heat makes particles move faster, increasing energy and causing expansion or phase changes.

Classifying Matter (Pure Substance, Solution, or Mechanical Mixture)

• Raisin bran cereal is a mechanical mixture because the different components are visible.
• Nitrogen is a pure substance (element).
• Sugar dissolved in water is a solution (homogeneous mixture).
• Iron is a pure substance (element).
• Stainless steel is a solution (homogeneous mixture/alloy).

Chemical vs. Physical Properties

• Physical properties describe a substance without changing its composition.
• Chemical properties describe a substance's ability to undergo chemical change.
• Examples of physical properties include the silver color of a computer, an ice cube floating on water, the transparency of glass, and ice melting at 0°C.
• Examples of chemical properties include reactivity with vinegar, flammability, and magnesium emitting white light when burned.

Chemical vs. Physical Changes

• A physical change does not create a new substance.
• A chemical change results in the formation of a new substance.
• Examples of physical changes include glass breaking, cutting hair, mixing oil and water, evaporating water, and mixing salt and water.
• Examples of chemical changes include bleaching hair, a rusting bicycle, and exploding fireworks.

Elements vs. Compounds

• A colorless gas that burns to produce CO₂ and H₂O is a compound, as it contains more than one element.
• A yellow solid with constant properties that cannot be broken down is an element.

Characteristic Physical Properties

• Characteristic physical properties do not change for a given substance and help identify it.
• Density (e.g., Water = 1 g/cm³), boiling point (e.g., Water = 100°C), and melting point (e.g., Ice = 0°C) are examples.

Density Calculations Using GRASP

• The GRASP method includes Given, Required, Analysis, Solution, Paraphrase.
• To calculate the density of nickel given a mass of 3.48 kg (3480 g) and a volume of 400 cm³, the density is 8.7 g/cm³.
• To calculate the density of a cube with a side length of 3.0 cm and a mass of 25 g, first find the volume (27 cm³), then calculate the density as 0.93 g/cm³.
• To calculate the volume of a rock with a density of 4 g/mL and a mass of 16 g, the volume is 4 mL.
• To calculate the mass of an irregularly shaped stone with an initial water volume of 20.0 mL, a final water volume of 30.2 mL, and a density of 2.45 g/mL, the mass is approximately 25 g.
• To determine the best wood for floating, calculate the density of each piece and choose the one with a density less than 1 g/cm³; Piece 2 with a density of 0.4 g/cm³ floats.

Atomic Theorists

• Democritus (400 BCE) proposed that matter is made of tiny, indivisible particles called "atomos."
• John Dalton (1803) developed the first atomic theory, stating that atoms are indivisible, atoms of the same element are identical, and atoms combine in fixed ratios to form compounds.
• J.J. Thomson (1897) discovered the electron using the cathode ray experiment and proposed the plum pudding model.
• Ernest Rutherford (1911) discovered the nucleus through the gold foil experiment.
• Niels Bohr (1913) proposed that electrons orbit the nucleus in specific energy levels, contributing the Bohr model.
• James Chadwick (1932) discovered the neutron, which explained atomic mass differences.

Location of Metals, Non-Metals & Metalloids on the Periodic Table

• Metals are located on the left and center of the periodic table (e.g., Sodium, Iron).
• Non-metals are located on the right side of the periodic table (e.g., Oxygen, Chlorine).
• Metalloids are located in a stair-step pattern between metals and non-metals (e.g., Silicon, Boron).

Properties of Metals vs. Non-Metals

• Metals conduct electricity, can be flattened by hammering (malleable), are often soft and shiny, and are located in the first column (alkali metals).
• Non-metals are often gases under normal conditions, are located in the upper-right corner, shatter when struck (brittle), and can be dull yellow powders.

The Four Main Chemical Families & Their Properties

• Alkali Metals (Group 1) are very reactive, soft, have low melting points, and possess 1 valence electron.
• Alkaline Earth Metals (Group 2) are less reactive than Group 1, have 2 valence electrons, and are harder metals.
• Halogens (Group 17) are highly reactive non-metals, form salts with metals, and have 7 valence electrons.
• Noble Gases (Group 18) are very stable (inert), have a full outer shell, and exhibit low reactivity.
• Reactivity increases down a group for metals and decreases down a group for non-metals.

Chemical Families of Elements

• Iodine (I) belongs to the Halogen family.
• Barium (Ba) belongs to the Alkaline Earth Metal family.
• Francium (Fr) belongs to the Alkali Metal family.
• Krypton (Kr) belongs to the Noble Gas family.

Unique Properties of Hydrogen

• Hydrogen is in Group 1 but is not an alkali metal.
• It is a non-metal and does not behave like other Group 1 elements.
• It can gain or lose an electron, making it highly versatile.

Maximum Electrons in Electron Shells

• The 1st shell can hold a maximum of 2 electrons.
• The 2nd shell can hold a maximum of 8 electrons.
• The 3rd shell can hold a maximum of 8 electrons (or up to 18 for larger atoms).

Common Traits Within Element Groups

• Elements in the same group have the same number of valence electrons.
• This determines their chemical properties and reactivity.
• Elements in the same group exhibit similar chemical behavior.

Atomic Structure & Diagrams

• To draw the Bohr-Rutherford/Lewis Dot Diagrams, first, determine the number of protons, electrons, and neutrons. Then, write the standard atomic notation and draw the Bohr-Rutherford diagram and Lewis dot diagram. Finally, draw diagrams for their ions.
• Lithium (Li) has 3 protons, 4 neutrons, 3 electrons, and an ion charge of +1.
• Sodium (Na) has 11 protons, 12 neutrons, 11 electrons, and an ion charge of +1.
• Oxygen (O) has 8 protons, 8 neutrons, 8 electrons, and an ion charge of -2.
• Magnesium (Mg) has 12 protons, 12 neutrons, 12 electrons, and an ion charge of +2.
• Boron (B) has 5 protons, 6 neutrons, 5 electrons, and an ion charge of +3.
• Chlorine (Cl) has 17 protons, 18 neutrons, 17 electrons, and an ion charge of -1.
• Neon (Ne) has 10 protons, 10 neutrons, 10 electrons, and an ion charge of 0.

Ions

• An ion is an atom that has gained or lost electrons to achieve a stable electron configuration.
• Metals (Groups 1-3) lose electrons to form positive ions (cations).
• Non-metals (Groups 15-17) gain electrons to form negative ions (anions).

Bohr-Rutherford Diagrams for Ions

• Iodine (I⁻) gains 1 electron to become an ion.
• Barium (Ba²⁺) loses 2 electrons to become an ion.
• Francium (Fr⁺) loses 1 electron to become an ion.
• Krypton (Kr) does not form an ion as it is a noble gas.

Finding Protons & Identifying an Element

• Protons = Mass number - Neutrons.
• An element with a mass number of 195 and 117 neutrons has 78 protons.
• Element with 78 protons = Platinum (Pt).

Reactivity of Metals

• Sodium (Na) is more reactive than Magnesium (Mg).
• Na has 1 valence electron (easier to lose), while Mg has 2 (more stable).

Reactivity of Non-Metals

• Chlorine (Cl) is more reactive than Oxygen (O).
• Cl is in Group 17 (halogens), which are highly reactive, while O is in Group 16.

Ionic or Molecular Compounds

• Potassium (K) and Fluorine (F) form an ionic compound (Metal + Non-Metal).
• Iron (Fe) and Chlorine (Cl) form an ionic compound (Metal + Non-Metal).
• Nitrogen (N) and Oxygen (O) form a molecular compound (Non-Metal + Non-Metal).
• Carbon (C) and Bromine (Br) form a molecular compound (Non-Metal + Non-Metal).

Diatomic Molecules

• Diatomic molecules consist of two atoms of the same element.
• The 7 diatomic molecules include H₂, N₂, O₂, F₂, Cl₂, Br₂, I₂.

Counting Atoms in Compounds

• List the elements present, then count the number of atoms of each element.
• NaHCO₃ (Sodium Bicarbonate): Na = 1, H = 1, C = 1, O = 3.
• CaCO₃ (Calcium Carbonate): Ca = 1, C = 1, O = 3.
• 5NaCl (Sodium Chloride): Na = 5, Cl = 5.
• 2KNO₃ (Potassium Nitrate): K = 2, N = 2, O = 6.
• Mg₃(PO₄)₂ (Magnesium Phosphate): Mg = 3, P = 2, O = 8.