All elements transfer the same number of electrons when bonding with other elements.
Electrons in metals can move freely among all the ions in the metal.
Some atoms bond by sharing electrons between the two atoms.
Water molecules have two opposite ends like the poles on a magnet.
1.1 Why Do Atoms Combine? I. Atomic Structure A. Atoms are mostly empty space B. Nucleus contains protons and neutrons C. Most of the volume is the electron cloud D. Electrons 1. NOT so much like the planets circling the sun 2. Planets do not have electrical charges 3. Planets move in predictable orbits
4. Cannot predict the exact position of an electron 5. Use a mathematical formula that predicts where it might be E. Element Structure 1. Specific number of protons, neutrons and electrons 2. Protons = electrons for a neutral atom
II. Electron Arrangement A. Electron Energy 1. Some electrons closer to the nucleus than others 2. Different areas where electrons are located = energy levels B. Number of Electrons 1. Each energy level can hold a max number of electrons 2. Farther from the nucleus = more electrons the energy level can hold
C. Energy Steps 1. Electrons closest to the nucleus have lowest energy and hardest to remove 2. Electrons farthest from nucleus have highest energy and easiest to remove 3. Formula: 2n^2 where n=energy level
III. The Periodic Table and Energy Levels A. For neutral atoms, # of protons = # of electrons B. Determine the number of electrons from atomic number IV. Electron Configurations A. Number of electrons increases from left to right on Periodic Table B. Atoms with full outer shells are stable C. Atoms with incomplete outer shells are unstable and will bond with other atoms
D. Elements in period 1: up to 2 electrons E. Elements in following period: up to 8 electrons V. Element Families (Groups) A. Elements in the same family have similar traits B. Noble Gasses 1. Eight electrons in outer shell 2. Stable (don’t react/form bonds) 3. Trait allows for many uses
C. Halogens 1. Need to gain 1 electron to be stable. 2. Fluorine a. Most reactive b. Outer energy level closest to nucleus. 3. Reactivity decreases down group 4. Outer energy levels are farther from nucleus
D. Alkali Metals 1. One electron in outer energy level 2. Need to lose 1 electron to be stable 3. Reactivity increases down the group 4. Outer energy levels farther from nucleus 5. Less energy needed to remove electron farther from nucleus.
VI. Electron Dot Diagrams A. Symbol for element surrounded by dots that represent electrons in outer energy level B. How to Write Them 1. Periodic table shows number of electrons in outer energy level. 2. Write element symbol 3. Start at top left and work clockwise by adding 1 dot to each side. 4. When more than 4, begin doubling up C. Using Dot Diagrams (on board)
1.2 How Elements Bond I. Ionic Bonds – Loss and Gain A. Sodium 1. Had 1 electron in outer shell 2. Makes it highly unstable 3. Reacts violently with water 4. Loses electron: stable outer shell B. Chlorine 1. Forms bonds the opposite way of sodium
2. Has 7 electrons in outer shell 3. Gains 1 electron to form stable octet C. Ions – A Question of Balance 1. Losing electron unbalances sodium’s charges 2. Becomes + charged 3. Chlorine becomes – charged 4. Ion: atom without a neutral 5. Become Na + and Cl -
D. Bond Formation 1. + charge is attracted to – charge 2. Attraction = ionic bond 3. Sodium chloride = table salt 4. Compound: a pure substance containing two or more elements that are chemically bonded. E. More Gains and Losses 1. Can elements gain/lose more than 1 electron?
2. Answer: yes! 3. Example: MgO and MgCl 2 II. Metallic Bonding – Pooling A. Metallic Bond: electrons move freely between atoms B. Many atoms share the same electrons C. Explains why metals are malleable, ductile, and conduct electricity. D. One pool of atoms and electrons; not individual molecules.
III. Covalent Bonds – Sharing A. The Covalent Bond 1. Share electrons only within the same molecule. 2. Molecule: neutral particle formed when atoms share electrons 3. Be able to explain the differences in ionic, metallic, and covalent bonds!
B. Double and Triple Bonds 1. Atoms share more than 1 pair of electrons 2. Examples: CO 2 and N 2 III. Polar and Non-Polar Molecules A. Some atoms attract electrons better than others B. Unequal sharing causes one side of the bond to be more – and the other more + C. Example: HCl
D. The Polar Water Molecule 1. Unequal sharing between H and O 2. Electrons more attracted to O 3. Molecules with out uneven charges are non-polar IV. Chemical Shorthand A. Symbols for Atoms 1. One or two letter abbreviations B. Symbols for Compounds 1. Use sub-scripts to indicate 2 or more
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