1. High School Physical Science
Play the song in the morning –not too loud with the visual words while everyone comes into the room. The Van de Graff generator is always such a positive experience!

2. Warm-up
An isotope of chlorine has 17 protons and 19 neutrons. What is the mass number for that isotope of chlorine? 2 17 19 36

3. What do we already know about the Periodic Table of Elements?
Opening KWL: Should already know: The three main groups of elements are metals, non-metals, and metalloids and their properties. The names of groups 1,2, 13, 14, 15, 16, 17, & 18. The number of valence shell electrons in each group. The period # equals the number of electron energy levels. Common symbols. That metals bond with non-metals to form ionic bonds. That non-metals form covalent bonds.

4. The amphoteric line, also known as the zig-zag line, divides metals and non-metals. T
There are only 6 metalloids so this table is not completely correct. Boron and astatine are non-metals. Aluminum is a metal. Metalloids: silicon, germanium, arsenic, antimony, tellurium, pollonium. The properties of metals versus nonmetals are given above.

5. Standard: Use IUPAC nomenclature for transition between chemical names and chemical formulas of
Binary ionic compounds (containing representative elements).
2. Binary covalent compounds (i.e.
carbon dioxide, carbon tetrachloride).
Discuss the terminology here: Binary means two. Ionic compounds are where ions are held together with ionic bonds. In ionic bonds there is unequal sharing of electrons.

6. Why do atoms bond?

7. This table shows the ions typically formed by groups 1,2, and 13 through 18. Groups 1,2, & 3 form cations. Groups form anions. The table also shows weak versus strong nuclear attraction for valence electrons.

8. Naming Type 1- Ionic Compounds
Vocabulary
Example:
Positive ion is called a “cation”
Negative ion is call an “anion”
Metal and non-metal
The cation is always named first and the anion second.
The cation takes the name of element
The anion takes the first part of the element name (the root) then adds –ide.
The net (overall) charge of a compound is always zero.
Sodium Chloride = Na+1 Cl-1
Rules
Compounds with two ions. When electrons are transferred (stolen), ions are formed. An anion gains electrons and is negatively charged (non-metals). The “n” in anion reminds me of negative. An anion becomes negatively charged. A cation loses electrons and is positively charged (metals) .

9. Naming Binary Ionic Compounds
Type 1 Ionic Compounds
Metals (and H) that form only one cation
Include H, Li, Na, K, Cs, Be, Mg, Ca, Ba, Al, Ag
Name the following compounds:
Rb2O SrI KS2
rubidium oxide
strontium iodide
potassium sulfide
There are three types of ionic compounds: Types 1, 2, and 3.

10. Naming Type 2 Binary Ionic Compounds
Rules
Practice: Give the systematic name for each of the following:
These metals can form two or more cations
Includes Cr, Cu, Fe, Sn, Pb, Co, Pb, Hg
Roman numeral specifies the charge of the ion
CuCl
HgO
Fe2O3
MnO2
PbCl4
copper(I) chloride
mercury(II) oxide
iron(III) oxide
manganese(IV) oxide
lead(IV) chloride
Some metals like copper can lose one or two electrons in a chemical reaction. By using a system with a roman numeral we can identify which ion of copper it is. "-ous" (lower has "o“ in it); "-ic" (higher has "i" in it)
Ion
Systematic Name
Older Name
Fe3+
iron (III)
ferric
Fe2+
iron (II)
ferrous
Cu2+
copper (II)
Cupric
Cu +
copper (I)
cuprous

11. Naming Type 3 Binary Compounds
Rules
Prefixes
Type 3 contain ONLY NONMETALS
The first element is named by its element name
The second element is named as though it is an anion with –ide
The prefix mono is never used for naming the 1st element. The others prefixes can be used for the 1st element.
Mono Di
Tri
Tetra
Penta
Hexa
Hepta
Octa
Examples:
BF3 = boron trifluoride
CO = carbon monoxide (not monocarbon monoxide)
N2O = dinitrogen monoxide (common name nitric oxide is used by some dentists as an anesthetic)
H2O= dihydrogen monoxide

12. Practice Naming Binary Compounds

13. Standard:
Apply the Law of Conservation of Matter by balancing the following types of chemical equations: synthesis, decomposition, single replacement, double replacement.

14. Law of Conservation of Mass
-In a chemical reaction, atoms are neither created nor destroyed. -Chemical equations must be balanced or have the same number and kinds of atoms on the product side of the equation as the reactant side.
The two main parts of a chemical equation are the reactants and products. The arrow indicates the direction of the change and is read as “yields” or “produces”. The coefficient tells you how many of this molecule or compound you have. This equation is balanced because you have 4 hydrogen and 2 oxygen in the reactants and 4 hydrogen and 2 oxygen in the product. This is an example of a synthesis reaction. The subscripts tell how many atoms are present in the molecule.

15. Writing Formulas for Compounds
Each atom present is represented by its element symbol
The number of each type of atom is indicated by a subscript written to the right of the element symbol
When only one atom of a given type is present, the subscript 1 is not written
1. How many atoms of each element are in SO3 ?
sulfer-1; oxygen-3
2. How many atoms of each element when adding a
coefficient such as 2S03?
sulfer-2; oxygen-6
With the coefficient of 2, there are a total of 2 sulfur and 6 oxygen atoms.

16. Balancing Chemical Equations
H2 + O2 → H2O
H= H=
O= O=
H2O2→ H2O + O2
Na + O2 →Na2O

17. 2Mg(s)  +  O2(g)  →    2MgO(s)
C(s)  +  O2(g)  →    CO2(g)
Ex: Hydrogen gas and oxygen gas burn to produce water.
2 H2 + O2 2 H2O and
sulfur trioxide reacts with
water to make sulfuric acid.
H2O + SO3 H2SO4
What would you see in a
"test tube" if you were witness to a synthesis reaction? You would see two different materials combine. A single new material appears.

18. Photosynthesis Reaction
This is an example of a balanced equation. The same type and number of atoms are in the 2 reactants as are in the 3 products. Photosynthesis synthesizes glucose.

19. Decomposition Reactions
Electrolysis of water is an example:
2 H2O → 2 H2 + O2
Another example is the decomposition of potassium chloride into potassium and chlorine gas:
2 KCl(s) → 2 K(s) + Cl2(g)

20. Replacement Reactions
a more active element takes the place of another element in a compound and sets the less active one free.
Replacement of a metal with a more active metal. EX. Fe(s)  +  CuSO4(aq)  →    FeSO4(aq)  +  Cu(s)
Replacement of hydrogen in water by an active metal. EX. 2Na(s)  +  2H2O(l)  →    2NaOH(aq)  +  H2(g)
EX. Mg(s)  +  H2O(g)  →    MgO(s)  +  H2(g)
Replacement of hydrogen in acids by active metals. EX. Zn(s)  +  2HCl(aq)  →    ZnCl2(aq)  +  H2(g)

21. Summary of Classifying Reactions

22. Test Questions
Which one of the following physical states is not used for H2O in a chemical equation?
a. (s) b. (l) c. (aq) d. (g)
What type of chemical reaction is this?
Cl2(g)  +  2NaBr(aq)  →    2NaCl(aq)  +  Br2(l)
3. Mg(s)+ O2(g) MgO(s)

23. Standard: Explain the periodic trends related to the types of ions formed.

26. The Periodic Table Atomic Radii Trends
Though there are some reversals in the trend (e.g., see Po in the bottom row), atoms generally get smaller as you go across the periodic table and larger as you go down any one column. Numbers are the radii in picometers (pm).

27. Increasing Ionization Energy
Ionization energy (IE) is the amount of energy required to remove an electron from an atom in the gas phase. IE is usually expressed in kJ/mol of atoms. It is always positive because the removal of an electron always requires that energy be put in (i.e., it is endothermic). IE also shows periodic trends. As you go down the periodic table, it becomes easier to remove an electron from an atom (i.e., IE decreases) because the valence electron is farther away from the nucleus. However, as you go across the periodic table and the electrons get drawn closer in, it takes more energy to remove an electron; as a result, IE increases:

28. Electronegativity is a measure of the attraction of an atom for electrons in a covalent bond.
Fluorine, the most reactive non-metal, is assigned the highest value since it has the greatest attraction for the electron being shared by the other element. Oxygen is also highly electronegative and has a strong attraction for electrons. Metals have low electronegativities since they have weak attraction for any shared electrons.
When two unlike atoms are convalently bonded, the shared electrons will be more strongly attracted to the atom of greater electronegativity. Such a bond is said to be polar. A polar bond results in the unequal sharing of the electrons in the bond.

30. Summary of Periodic Table Trends
Moving Left → Right
Atomic Radius Decreases
Ionization Energy Increases
Electronegativity Increases
Moving Top → Bottom
Atomic Radius Increases
Ionization Energy Decreases
Electronegativity Decreases

31. Covalent bonds are formed by diatomic and polyatomic molecules due to equal electronegativities.

33. Test Questions
Which combination of atoms can form a polar covalent bond?
a. H and H b. H and Br c. N and N d. Na and Br
The electronegativity difference between H (2.1) and Br (3.0) is 0.9. If the difference is less than 1.7 (but not zero, usually more than 0.4), a polar covalent bond is formed. A nonpolar covalent bond is formed when the difference is close to zero (usually up to 0.4). If the difference is more than 1.7, an ionic bond is formed.
A strontium atom differs from a strontium ion in that the atom has a greater
a. # of electrons b. # of protons c. atomic # d. mass #
Which bond has the greatest ionic character (difference in electronegativity)?
a. H—Cl b. H—F c. H—0 d. H--N
H has an electronegativity of 2.1 and F has an electronegativity of 4.0. Subtract the two values and the difference is 1.9. Ionic bonds have differences of 1.7 or greater.