Chem Deck Flashcards
Chem test (43 cards)
Metals, Non-Metals, Semi-Metals
Metals: Conduct electricity, malleable, shiny, lose electrons to form positive ions (cations).
Non-Metals: Poor conductors, brittle, gain electrons to form negative ions (anions).
Semi-Metals: Have properties of both metals and non-metals (e.g., silicon).
Groups & Periods
Groups (Columns): Elements in the same group share similar properties (e.g., Group 1: alkali metals, Group 17: halogens).
Periods (Rows): Elements in the same period have the same number of electron shells, but their properties change across the period.
Subatomic Particles
Protons: Positive charge, mass of 1, found in the nucleus.
Neutrons: No charge, mass of 1, found in the nucleus.
Electrons: Negative charge, no mass, orbit around the nucleus.
Atomic Number, Mass, & Neutrons
Atomic Number: Number of protons, defines the element.
Mass Number: Protons + Neutrons.
Neutrons: Mass - Atomic Number.
Example: Oxygen has a mass number of 16 and 8 protons, so it has 8 neutrons.
Isotopes
Same element, different mass due to different numbers of neutrons. Example: Hydrogen-1 (1 proton, 0 neutrons), Hydrogen-2 (1 proton, 1 neutron), Hydrogen-3 (1 proton, 2 neutrons).
Average atomic mass on the periodic table is the weighted average of all isotopes.
Bohr-Rutherford Diagram
Shows protons and neutrons in the nucleus and electrons in shells.
Example: Fluorine has 9 protons, 9 electrons, and 10 neutrons.
Electron Shells (Max # of Electrons per Shell)
1st shell = 2 electrons
2nd shell = 8 electrons
3rd shell = 8 electrons
4th shell = 2 electrons (for elements like calcium)
Valence Electrons
Group 1 = 1 valence electron, Group 2 = 2 valence electrons, etc.
Lewis Dot Diagrams
Draw the element symbol and place valence electrons (dots) around it.
Example: Fluorine (Group 17) will have 7 dots (one for each valence electron).
Ions
Cation: Positive charge, loses electrons (e.g., Na+).
Anion: Negative charge, gains electrons (e.g., O2-).
How Ions Form
Metals lose electrons to form cations (positive charge).
Non-metals gain electrons to form anions (negative charge).
Octet Rule
Elements aim to have 8 electrons in their outer shell (except for hydrogen and helium, which need 2 electrons).
Ionic Compounds
Formed when a metal transfers electrons to a non-metal.
Example: Sodium (Na) + Chlorine (Cl) → NaCl.
Ionic Bonding:
Metals lose electrons to non-metals, forming ions that are oppositely charged and attract each other.
Molecular Compounds
Formed when two non-metals share electrons.
Example: Oxygen (O) + Oxygen (O) → O2 (double bond).
Lewis Diagrams for Ionic Compounds
Example: Magnesium (Mg) + Chlorine (Cl) → MgCl2.
Magnesium loses 2 electrons (Mg2+), chlorine gains 1 electron (Cl-).
Write the formula by “criss-crossing” charges: MgCl2.
Ionic Compound Naming
Naming ionic compounds: Metal + Non-metal (change the non-metal’s ending to “ide”).
Example: NaCl = Sodium Chloride.
Molecular Compound Naming
Use prefixes:
1 = mono, 2 = di, 3 = tri, etc.
Example: CO2 = Carbon Dioxide.
Molecular Compound Naming (numbers)
“Mono-” indicates one, “di-” indicates two, “tri-” is three, “tetra-” is four, “penta-” is five, “hexa-” is six, “hepta-” is seven, “octo-” is eight, “nona-” is nine, and “deca” is ten. If there is only one of the first element, you can drop the prefix.
Atomic Size
Down a Group: Atoms get bigger (more shells).
Across a Period: Atoms get smaller (more protons attract electrons).
Reactivity
Metals: As you go down a group, reactivity increases because the attraction to valence electrons weakens (more shells).
Non-metals: As you go up a group, reactivity increases because the atom is smaller and the nucleus attracts electrons more strongly.
Metals
Lose electrons to form cations (positive charge).
React with water (alkali metals are highly reactive).
Non-Metals
Gain electrons to form anions (negative charge).
Stable like noble gases (Group 18) when they achieve a full valence shell.
Isotopes
Same element, different number of neutrons (affects mass number).
