Chemistry Chapter 5 - Atomic Structure and Bonds

A pure substance that cannot be broken down into simpler substances by chemical means.

Two or more elements chemically combined in a fixed ratio.

A mixture where the components are evenly distributed throughout the mixture.

A mixture where the components are not evenly distributed throughout the mixture.

A homogeneous mixture where a solute dissolves in a solvent.

The substance that dissolves in a solvent to form a solution.

The substance that dissolves the solute to form a solution.

A heterogeneous mixture where particles are suspended in a liquid, but settle over time.

A technique used to separate mixtures based on the different affinities of the components for the stationary and mobile phases.

Chromatography can separate mixtures where components have different polarities or affinities for the stationary phase.

An endothermic process is a process that absorbs heat from the surroundings, resulting in a decrease in the temperature of the surroundings.

An exothermic process is a process that releases heat into the surroundings, resulting in an increase in the temperature of the surroundings.

Alkali metals are highly reactive, soft, silvery metals that react readily with water to form strong bases.

Alkaline earth metals are reactive, silvery-white metals that are less reactive than alkali metals but still readily react with water to form strong bases.

Halogens are highly reactive, nonmetals that exist as diatomic molecules and readily form ionic compounds with metals.

Noble gases are unreactive, colorless gases that exist as single atoms and have a full outer shell of electrons, making them very stable.

The minimum energy required to remove one electron from a gaseous atom in its ground electronic state.

A measure of the relative ability of an atom in a molecule to attract electrons to itself.

The ability of a substance to conduct electricity when dissolved in water or melted.

A chemical bond formed by the sharing of electrons between two nonmetals.

The ability of a substance to conduct electricity in its molten or dissolved state.

The process of a metal atom losing an electron to form a positive ion.

As the number of protons in the nucleus increases, the attractive force between the nucleus and electrons also increases, resulting in a smaller atomic radius.

An ionic bond is formed through the electrostatic attraction between a positively charged ion (cation) and a negatively charged ion (anion)

The ability of a substance to be stretched into a wire without breaking.

The ability of a substance to be hammered or pressed into different shapes without breaking.

In metals, valence electrons are not bound to any specific atom and can move freely throughout the metal lattice.

The ability of a substance to conduct electricity.

A bond where electrons are shared unequally between two atoms, resulting in a partial positive and partial negative charge.

Intermolecular forces (IMFs) are attractive forces that exist between molecules. Dipole-dipole interactions occur between polar molecules, hydrogen bonding occurs between molecules with H bonded to F, O, or N, and London dispersion forces occur between all molecules, but are stronger in larger molecules.

The substance with the stronger intermolecular forces will have a higher boiling point.

The sum of the atomic masses of all the atoms in a molecule.

The empirical formula represents the simplest whole-number ratio of atoms in a compound. To find it, you need to convert the percentages of each element to grams, then to moles. Divide each mole value by the smallest mole value to get the ratio. For example, let's say we have a compound with 38.8% carbon, 16.2% hydrogen, and 45.1% nitrogen. Assuming a 100 g sample, we get 38.8 g of carbon, 16.2 g of hydrogen, and 45.1 g of nitrogen. Convert these masses to moles: 38.8 g C / 12.01 g/mol = 3.23 mol C, 16.2 g H / 1.01 g/mol = 16.0 mol H, and 45.1 g N / 14.01 g/mol = 3.22 mol N. Divide each mole value by the smallest mole value (3.22 mol): 3.23 mol C / 3.22 mol = 1.00, 16.0 mol H / 3.22 mol = 5.00, and 3.22 mol N / 3.22 mol = 1.00. Therefore, the empirical formula is CH5N.

The molecular formula represents the actual number of atoms of each element in a molecule. To find it, you need to know the empirical formula and the molar mass of the compound. Determine the empirical formula mass, then divide the molar mass by the empirical formula mass. This result tells you how many times the empirical formula needs to be multiplied to get the molecular formula. For example, if we have a compound with an empirical formula of CH5N and a molar mass of 124.10 g/mol, the empirical formula mass is 12.01 + (5 x 1.01) + 14.01 = 31.06 g/mol. Divide the molar mass by the empirical formula mass: 124.10 g/mol / 31.06 g/mol = 4. Therefore, the molecular formula is 4 x CH5N = C4H20N4.

The empirical formula of an ionic compound represents the simplest whole-number ratio of ions in the compound. To find it, convert the masses of each element to moles, then divide each mole value by the smallest mole value to get the ratio. For example, let's say we have a compound with 34.95 g of iron and 15.05 g of oxygen. Convert these masses to moles: 34.95 g Fe / 55.845 g/mol = 0.627 mol Fe, 15.05 g O / 16.00 g/mol = 0.941 mol O. Divide each mole value by the smallest mole value (0.627 mol): 0.627 mol Fe / 0.627 mol = 1.00, 0.941 mol O / 0.627 mol = 1.50. Since we need whole numbers, multiply both by 2: Fe2O3. Therefore, the empirical formula is Fe2O3.

The molecular formula of a compound represents the actual number of atoms of each element in a molecule. You can find it by knowing the empirical formula and the molecular weight of the compound. Firstly, calculate the empirical formula mass. Then, divide the molecular weight of the compound by the empirical formula mass. This result will tell you how many times the empirical formula has to be multiplied to obtain the molecular formula. For example, if galactose has a molecular weight of 180.16 g/mol and an empirical formula of CH2O, the empirical formula mass = 12.01 + (2 x 1.01) + 16.00 = 30.03 g/mol. Dividing the molecular weight (180.16 g/mol) by the empirical formula mass (30.03 g/mol) gives us approximately 6. This means the molecular formula is 6 x CH2O = C6H12O6.

Physical properties are characteristics that can be observed or measured without changing the chemical composition of the substance. Examples include color, odor, taste, density, melting point, boiling point, solubility, and state of matter. These properties can be used to identify an unknown compound by comparing them to known properties of substances.

Percent composition refers to the percentage by mass of each element in a compound. To find the percent composition of a specific element, you need to divide the mass of that element in the compound by the total mass of the compound, and then multiply by 100%. For example, in acetylsalicylic acid (C9H8O4), the percent composition of oxygen is calculated as follows: 1. Calculate the mass of oxygen: 4 x 16.00 g/mol = 64.00 g/mol. 2. Calculate the total mass of the compound: (9 x 12.01) + (8 x 1.01) + (4 x 16.00) = 180.15 g/mol. 3. Divide the mass of oxygen by the total mass of the compound: 64.00 g/mol / 180.15 g/mol = 0.355. 4. Multiply by 100% to express it as a percentage: 0.355 x 100% = 35.5%. Therefore, the percent composition of oxygen in acetylsalicylic acid is 35.5%.

Particle diagrams are simplified representations of the arrangements of atoms and/or molecules in a substance. They can be used to visualize the difference between pure substances, mixtures, and different states of matter. For example, a particle diagram of pure gaseous CO2 would show individual CO2 molecules separated from each other, while a particle diagram of a mixture of gaseous CO2 and O2 would show both types of molecules mixed together. Remember that in a solid, particles are packed closely together in a regular arrangement, while in a liquid, particles are closer together but can move past each other, and in a gas, particles are far apart and move freely.