1. Chemistry Q1 Amazing Benchmark Review

2. Example 1: Standard 1a: Know how to relate the position of an element in the periodic table to its atomic number and atomic mass.

3. Example 1:

4. Example 2: Standard 1b: Know how to use the periodic table to identify metals, semimetals (metalloids), and nonmetals.

5. Example 3: Metals, Nonmetals, and Metalloids Metalloids

6. Example 3: Standard 1c: Know how to use the periodic table to identify alkali metals, alkaline Earth metals, and transition metals, halogens, noble gases, trends in ionization energy, electro-negativity, and the relative sizes of ions and atoms.

7. Example 2:

8. Atomic Size Size tends to increase down a column. Size tends to decrease across a row. As principal energy level increases, average distance of electrons from the nucleus also increases.

9. –Metals have low ionization energies –Tends to decrease down a group or column –Tends to increase across a period or row Ionization is the energy required to remove an electron from an individual atom in the gas phase.

10. Electro-negativity – the relative ability of an atom in a molecule to attract shared electrons to itself – Increases from left to right across a period – Decreases down a group of representative elements

11. Relative sizes of ions and atoms Anions are always LARGER than the parent atom - gain electron(s) Cations are always SMALLER than the parent atom - lost electron(s)

12. Example 4: Standard 1d: Know how to use the periodic table to determine the number of electrons available for bonding.

13. Example 4: Classifying Electrons Valence electrons – electrons in the outermost (highest) principal energy level of an atom Core electrons – inner electrons Elements with the same valence electron arrangement show very similar chemical behavior.

14. Example 4: Octet and Duet Rule In determining the number of electrons available for bonding, we include only the valence electrons (VE) Most important requirement for the formation of a stable compound is – Atoms achieve noble gas electron configuration using the octet rule (share 8e-) and duet rule (share 2e-).

15. Example 5: Standard 1e: Know the nucleus of the atom is much smaller than the atom yet contains most of its mass.

16. Example 5:

17. Example 6: Standard 2a: Know atoms combine to form molecules by sharing electrons to form covalent or metallic bonds or by exchanging electrons to form ionic bonds.

18. Chemical Bonds A covalent bond results when electrons are shared by two or more nuclei Covalent Bonding

19. Example 6: Cations (lose electron(s))

20. Example 6: Anion (gain electron(s))

21. Example 6: With ionic bonds, the charges on the anions and cations in the compound must equal to zero.

22. Example 7: Standard 2b: Know chemical bonds between atoms in molecules such as H2, CH4, NH3, H2CCH2, N2, Cl2, and many large biological molecules are covalently attracted to each other. Note: A covalent bond results when electrons are shared by two or more nuclei

23. Example 7 Covalent bonds

24. Example 8: Standard 2c: Know salt crystals, such as NaCl, are repeating patterns of positive and negative ions held together by electrostatic attraction. Note: Bonding forms Lattice Structure

25. Example 8: Lattice Structure Ions are packed together to maximize the attractions between ions. For example, Lithium Fluoride (LiF)

26. Example 9: Standard 2e: Know how to draw Lewis dot structures.

27. In writing Lewis structures, we include only the valence electrons (VE).