Chem notes 7

Exam Information

• The test will be online next week, available on your own time. Regular class attendance is still expected.
• There will be a symbol quiz focusing on the symbols and names from the first slide of Chapter Four. No need to know positions or numbers.

Atomic Models

• Democritus and Dalton: Proposed the solid sphere model of the atom, resembling a pool ball.
• J.J. Thompson:
  • Discovered the electron.
  • Proposed the plum pudding model, where electrons are scattered throughout a positive background.
• Rutherford:
  • Conducted the gold foil experiment, observing positive particles ricocheting back.
  • Proposed the nuclear model with a central nucleus containing protons.
  • Established that atoms are neutral, with the number of electrons equaling the number of protons.
  • Determined that atoms are mostly empty space: if an atom were the size of a stadium, the nucleus would be the size of a pea, and the electrons would be spread throughout the stadium.
  • Electron-electron repulsion prevents objects from passing through each other.
• Bohr (Neils Bohr):
  • Developed the solar system model (planetary model), where the nucleus is the sun, and electrons orbit like planets.
  • Postulated that electrons orbit the nucleus at specific distances and energy levels called orbits.
  • Each orbit has a specific distance from the nucleus and a specific energy.
  • Electrons can jump to higher energy levels when energy is applied and release energy as light when they fall back down.
  • The Bohr model is useful for explanations but has limitations.
• Schrodinger:
  • Proposed the electron cloud model or quantum mechanical model.
  • Replaced orbits with orbitals, which are 3D regions of space where electrons are likely to be found.
  • Orbitals are probability maps; the exact location of an electron is not predictable.
  • Introduced shapes of orbitals: S (sphere) and P (dumbbell).

Orbitals

• S Orbital:
  • Spherical shape in 3D space.
  • Holds up to two electrons.
• P Orbitals:
  • Dumbbell shape.
  • More complex shapes exist for D and F orbitals.

Electron Configurations

• Electron configurations describe where all the electrons go within an atom (to be revisited in Chapter Nine).
• Example: 1s^2 2s^2 2p^6 3s^2 3p^6
• Each orbital holds two electrons, superimposed with the nucleus in the center.
• The electron cloud model explains more phenomena regarding light than the Bohr model.

Neutron

• Chadwick: Discovered the neutron in the nucleus.
• Neutrons are neutral (no charge).
• Neutrons act as a "glue" to hold the protons together in the nucleus, counteracting proton-proton repulsion.
• The proton-to-neutron ratio determines the stability of an atom; unstable atoms are radioactive.
• Lighter elements (e.g., carbon, boron) are often stable with an equal number of protons and neutrons.
• Heavier elements (e.g., gold, mercury) require more neutrons to stabilize the protons.
• Example: Zirconium (Zr) with 40 protons is stabilized by approximately 45 neutrons.
• Excessive neutrons lead to radioactive decay, such as beta particle emission.

Subatomic Particles Chart

• Atomic Mass Unit (AMU) is used to measure individual atoms or molecules.
• One AMU is defined as one-twelfth of the mass of a carbon-12 atom.
• Carbon-12 has six protons and six neutrons.
• Carbon Symbol: ^{12}_{6}C
  • 6 is the atomic number (number of protons).
  • 12 is the mass number (sum of protons and neutrons).

Mass Comparisons

• If a proton were the weight of a baseball, an electron would be the weight of a grain of rice.

Element Symbols

• Symbols indicate the name of the element and are mostly English, with some Greek and Latin origins.
• Example: Na for sodium comes from natrium.
• If multiple elements start with the same letter, a second or third letter is added to differentiate (e.g., Ca for calcium, Ce for cerium, Cs for cesium).

Atomic Number

• The atomic number is the number of protons in an atom.
• It is the whole number on the periodic table.
• Example: Magnesium (Mg) is number 12 on the periodic table, so it has 12 protons.

Historical Context

• Alchemists attempted to transmute lead into gold, which is impossible without modern technology like particle accelerators because lead and gold have different numbers of protons.

Periodic Table Organization

• Dimitri Mendeleev:
  • Considered the father of the periodic table.
  • Organized elements by relative mass and observed recurring patterns in their properties.
  • Arranged 65 elements, grouping them into columns (groups or families) with similar properties.
  • Rows are called periods, and there is no pattern of reactivity from left to right in a period.
  • Predicted the properties of undiscovered elements based on their position in the table.

Periodic Table Annotations

• Numbering: Use the A-B system to indicate the number of outside electrons, which is helpful for understanding regular chemical reactions.
• Group 1A: Place above hydrogen
• Group numbering of 2A and then 3A through 8A.
• Staircase: Draw a darker line to represent the staircase, which separates metals from non-metals (left and right, respectively).
• Aluminum is an exception; it touches the staircase but is a metal.
• Main Group Metals: Groups 1A and 2A and under the staircase.
• Transition Metals: The B groups.
• Inner Transition Metals (Rare Earth Metals): The two rows at the bottom, also known as:
  • Lanthanides (top row, starting with La).
  • Actinides (bottom row, starting with Ac).

Properties of Metals

• Good conductors of heat and electricity due to free electrons.
• Malleable: Can be pounded into different shapes.
• Ductile: Can be drawn into wires.
• Shiny.
• Lose electrons.

Properties of Non-Metals

• Bad conductors of heat and electricity (insulators).
• Dull and brittle (solids).
• Exist as solids, liquids, and gases.
• Gain electrons.

States of Matter

• Most metals are solid at room temperature; mercury (Hg) is the only liquid metal at room temperature.
• Non-metals exist in all states of matter (solid, liquid, gas).

Metalloids (Semimetals)

• Elements touching the staircase (except aluminum).
• Have properties in between metals and non-metals.
• Example: Silicon (Si), a semiconductor, conducts electricity well when warm but poorly when cold.

Group 1A: Alkali Metals

• Most reactive metals with one outer electron.
• Readily lose one electron to achieve stability, resembling a noble gas.
• React violently in water.
• Going down the group, atoms get larger, making it easier to lose the outer electron.

Group 2A: Alkaline Earth Metals

• Have two outer electrons.
• Less reactive than alkali metals.
• Always found as part of compounds in nature (e.g., calcium carbonate).