Unit 1: Atomic Structure and Properties

21 terms. Say what each one means, then open the card to check.

• Mole

  The amount of a substance that holds 6.022 × 10²³ of its particles, so chemists can count atoms and molecules by weighing them.

  Example. A balloon holding 2.00 mol of helium contains 2.00 × 6.022 × 10²³ = 1.20 × 10²⁴ helium atoms.

• Avogadro's number

  The number of particles in one mole, 6.022 × 10²³ per mole. It converts a count of particles into moles and back.

  Example. A sample of neon that holds 3.011 × 10²³ atoms holds 3.011 × 10²³ ÷ 6.022 × 10²³ = 0.5000 mol of neon.

• Molar mass

  The mass in grams of one mole of a substance, in g/mol. Add the atomic masses of every atom in the formula.

  Example. For sulfur dioxide, SO₂: 32.06 + 2(16.00) = 64.06 g/mol.

• Percent composition by mass

  The share of a compound's mass that comes from each element, found from the formula or from measured masses.

  Example. In water, hydrogen supplies 2(1.008) = 2.016 g of every 18.02 g, so water is 11.2% hydrogen by mass.

• Empirical formula

  The simplest whole-number ratio of the atoms of each element in a compound.

  Example. Glucose is C₆H₁₂O₆, and dividing every subscript by 6 gives its empirical formula, CH₂O.

• Molecular formula

  The actual number of atoms of each element in one molecule. It is a whole-number multiple of the empirical formula.

  Example. A compound with empirical formula CH (13.02 g/mol) and molar mass 78.11 g/mol has 78.11 ÷ 13.02 ≈ 6 units, so it is C₆H₆.

• Law of definite proportions

  A pure compound always contains its elements in the same ratio by mass, wherever the sample came from.

  Example. Pure water from a lake or from a lab is always 16.00 ÷ 18.02 = 88.8% oxygen by mass.

• Isotope

  Atoms of the same element, so the same number of protons, that have different numbers of neutrons and so different masses.

  Example. Carbon-12 and carbon-13 both have 6 protons; carbon-12 has 6 neutrons and carbon-13 has 7.

• Mass spectrum

  A graph of relative abundance against mass. For an element, each peak is one isotope and its height shows how common that isotope is.

  Example. A sample of bromine gives two peaks of almost equal height, at about 79 and 81 amu, so its two isotopes are about equally common.

• Average atomic mass

  The weighted average of an element's isotope masses, each multiplied by its fraction of the atoms. It is the mass shown on the periodic table.

  Example. If 90% of an element's atoms have mass 40 amu and 10% have mass 42 amu, the average is 0.90(40) + 0.10(42) = 40.2 amu.

• Electron configuration

  A list of the occupied subshells of an atom or ion with the number of electrons in each, filled from the lowest energy up.

  Example. Aluminum has 13 electrons: 1s² 2s² 2p⁶ 3s² 3p¹, or [Ne] 3s² 3p¹ in shorthand.

• Aufbau principle

  In the ground state, electrons fill the lowest-energy subshells first. The 4s subshell fills before 3d.

  Example. Potassium's nineteenth electron goes into 4s, giving [Ar] 4s¹, not into 3d.

• Valence electrons

  The electrons in an atom's outermost shell, the highest n. They are the farthest from the nucleus and take part in bonding.

  Example. Phosphorus is [Ne] 3s² 3p³, so it has 2 + 3 = 5 valence electrons in its third shell.

• Shielding

  Inner (core) electrons repel the outer electrons, so the outer electrons feel only part of the nuclear charge.

  Example. Lithium's 2s electron is shielded by its 2 core electrons, so it feels far less than the full pull of 3 protons.

• Effective nuclear charge

  The net positive charge a valence electron feels, roughly the number of protons minus the number of core electrons.

  Example. Sulfur has 16 protons and 10 core electrons, so its valence electrons feel a charge of about +6.

• Coulomb's law

  The force between two charges grows with the size of the charges and weakens with distance: F is proportional to q₁q₂ ÷ r².

  Example. If an electron were twice as far from the same nuclear charge, the attraction would be one quarter as strong.

• Photoelectron spectroscopy (PES)

  A method that knocks electrons out of atoms with high-energy light and records how much energy each electron needed. Each peak is one subshell.

  Example. Neon's spectrum has three peaks, for 1s, 2s and 2p, with heights in the ratio 2 : 2 : 6.

• Binding energy

  The energy needed to remove an electron from a given subshell. The closer and less shielded the electron, the higher it is.

  Example. In magnesium, the 1s electrons need far more energy to remove than the 3s electrons, so the 1s peak sits far from the 3s peak.

• Ionization energy

  The energy needed to remove the most loosely held electron from a gaseous atom or ion. Each further electron takes more energy.

  Example. Removing one electron from each atom in a mole of gaseous lithium takes about 520 kJ.

• Atomic radius

  The size of an atom. It shrinks across a period as nuclear charge rises, and grows down a group as electrons fill higher shells.

  Example. Across period 2, the atoms shrink steadily from lithium to fluorine, even though fluorine has more electrons.

• Electronegativity

  How strongly an atom attracts the shared electrons in a bond. It rises across a period and falls down a group.

  Example. In an H–F bond the shared electrons sit closer to fluorine, the most electronegative element.