1. Unit 4 – Compounds and Bonding

2. What Do You Remember? 1) Define Compound
Substance formed by 2 or more elements in which proportions are always the same
2) Give two example of a compound
H20, NaCl, CO2
3) Define Element
Substance with atoms that are all alike
4) Give two example of an element
Any from the periodic table

3. 6) Give an example of a mixture
5) Define Mixture
A material made up of 2 or more substances that can be easily separated by physical means
6) Give an example of a mixture
Soda, salad dressing, soup mixes, granite, cheese
7) What are 2 physical properties of iron?
Malleable, luster, ductile, good conductor of heat and electricity

4. 8) What are 2 physical properties of sulfur?
- solid at room temp, brittle, no luster, not malleable
9) How is iron sulfide different from the 2 elements that make it up?
- when you combine iron and sulfur it smokes
10) When iron sulfide is formed, is it a physical or a chemical change.
- chemical change (making a new substance)

5. a) Can be separated by physical means?
11) If you mixed the iron and sulfur together, would that be a chemical or physical change?
Physical – can separate by using a magnet
12) Mixture or Compound?
a) Can be separated by physical means?
mixture
b) No definite chemical composition
c) Maintains same properties as original compounds

6. 13) Draw a Bohr Diagram of Nitrogen

7. 14) Mixture, Compound, or Element
a) Salt Water
- mixture
b) Salt
- compound
c) Phosphorus
- element
d) Hydrochloric Acid
e) Einsteinium
f) Sugar
g) Hydrogen Peroxide
h) Carbon Monoxide

8. Stability in Bonding
Sodium (Na) and Chlorine (Cl) together can make up sodium chloride (NaCl)
Compounds such as sodium chloride have properties that aren’t like any of those of the individual elements
A chemical formula - tells what elements a compound contains and the exact number of atoms of each element
NaCl – 1 molecule of sodium and 1 molecule of Chlorine

9. Some Familiar Compounds
Water – H2O
Salt - NaCl
Sucrose – C12H22O11
Glucose – C6H12O6
Hydrochloric Acid – Stomach Acid – HCl
Sulfuric Acid – Battery Acid – H2SO4
Dinitrogen oxide – Laughing Gas – N2O

10. Most of the elements can combine with other elements
The 6 noble gases in Group 18 rarely form compounds because the atoms are unusually stable

11. Electron Dot Diagrams
Only show the electrons in the outer energy level of an atom – Valence Electrons
Don’t have to show energy levels

12. VALENCE ELECTRONS Elements in group 1 have 1 outer electron
Elements in group 2 have 2 outer electrons
Skip transitional metals (the energy levels are similar to groups 1 and 2)
Elements in group 13 have 3 outer electrons
Elements in group 14 have 4 outer electrons, and so on

13. Steps to Drawing Dot Diagrams:
Write chemical symbol
Find how many valence electrons are in element
Draw dots around symbol

15. Let’s try drawing some of these…K Mg Al Na Cl

16. Atoms with partially stable outer energy levels can lose, gain, or share electrons to obtain a stable outer energy level.
Note that there is 1 electron in the outer level of Na and 7 in Cl.
When they combine, sodium loses one electron and chlorine gains one
Hydrogen and oxygen share electrons to become more stable in the compound water

17. A chemical bond is the force that holds atoms together in a compound.

18. Ionic and Covalent Bonding
Why do most elements form compounds?
Become more stable – atomic stability
More resistance to change
Elements are most stable when they have 8 electrons on their outermost energy level
MAGIC # is 8!
Atoms are happy  when their outer energy levels are full!

19. Valence Electrons – the electrons that are involved in forming chemical bonds (outer energy level)
Correspond to the groups; elements in the same group have the same number of valence electrons
Except He (only has 2 electrons) and transitional metals vary

20. Examples:

21. How many valence electrons do the following atoms have?
Na – 1
Ar – 8
F – 7
In general, an atom will form a chemical bond with another atom so that both atoms will have 8 electrons in their outer energy level. (Exception: Noble Gases)

22. Two main types of chemical bonds:
Ionic
Covalent

23. Ionic Bonds (IT – Ionic Transfers electrons )
Force of attraction between opposite charges of ions
Transfer electrons
Non-metal and metal
Ex. NaCl
High boiling point, high melting point, conducts electricity
Crystalline solids

24. Covalent Bonds (CS – Covalent Shares electrons)
Attraction that forms between atoms when they share electrons
Non-metal and non-metal
Ex. H2O
Low boiling point, low melting point, doesn’t conduct electricity
Molecules

25. Identify the element, number of valence electrons and would the element develop a covalent bond, ionic bond, both or neither.

26. Organic Compounds
Organic compounds – most of the compounds that contain the element CARBON
Organic compounds make up all living things or things that have lived

27. 90% of all compounds are organic!

28. Compounds such as carbon dioxide and the carbonates are considered inorganic.
Inorganic compounds are mostly compounds not containing carbon

29. Why can carbon form so many compounds?
4 electrons in outer shell
4 covalent bonds
Can bond with many elements (especially hydrogen and oxygen)
Can link with other carbon atoms in different arrangements C

30. To determine molecular structure of organic compounds use the following formula:
CnH2n + 2

31. Examples of organic molecules
Methane gas (CH4) H H C H H

32. Propane gas molecule (C3H8)

33. Hydrocarbons – a compound made up of only carbon and hydrogen atoms
Hydrocarbons are currently the main source of the world’s electrical energy and heat sources (such as home heating) because of the energy produced when burnt
Ex. Methane, propane, butane, octane

34. Writing Formulas: Criss-Cross Method
1) Write down chemical symbols
2) List oxidation numbers in the upper right hand corner of each symbol
Oxidation number – tells you how many electrons an atom has gained, lost, or shared to become stable
To find the oxidation number take the – valence electrons. (for groups 13-18)

35. Oxidation Number of Groups:
Group 18 = 0 (remember…they don’t combine anyway)

36. 4) Criss-cross the oxidation numbers to the new location
3) Next to the number in the right hand corner, write down + for metals, and – for non-metals.
The element with the + oxidation number is written first
The element with the – oxidation number is written second
4) Criss-cross the oxidation numbers to the new location
5) Drop the + or – and rewrite the formula.

37. Example:

38. Examples:
Ag + S →
Li + F →
Ba + Cl →
Ca2 + B-5 →
Fe+2 + Ar →

39. Chemical Reactions What is a chemical reaction?
When substances go through a chemical change and form new substances with different properties.
What are some examples of chemical reactions?
Rusting of iron
Souring of milk
Burning of paper
Cooking an egg

40. 2 Parts of a chemical reaction:
Reactants - substances that enter into a reaction
Na + Cl  NaCl
Products - substances produced by a reaction
Example: What are the reactants and products in the rusting of iron?
Reactants Iron, water and oxygen
Products Rust (Iron Oxide)

41. Rusting is a chemical reaction
Rusting is a chemical reaction. The reactants are iron and oxygen and the product is iron oxide (rust).
For rusting to occur, the iron atoms must lose some of their electrons (negative particles that travel around the nucleus of an atom), and the oxygen must gain them.
Water is needed for rusting to occur because the electrons that the iron loses travel through the water.

42. How does a chemical reaction occur?
An atom will try to get 8 valence electrons by breaking existing bonds and rearranged to form new bonds.
A convenient way to express a chemical reaction is by writing a chemical equation:
4Fe O  Fe2O3
coefficient reactants “yields” products

43. http://www.youtube.com/watch?v=dExpJAECSL8 Law of Conservation of Mass
Matter cannot be created or destroyed during a chemical reaction (it can only be rearranged)
The total mass of the reactants EQUALS the total mass of the products.
Example:
2H O2  2H2O
If the total of 5 grams of Hydrogen and Oxygen gases were combined, then ______ of water would be obtained.
Answer: 5 grams

44. Balanced Chemical Equations
A chemical equation must show that atoms are NOT created or destroyed.
Must have the same number of atoms on both sides
Cannot add subscripts – use coefficients to balance equations

45. Is this equation balanced?
H2 + O2 → H2O
No there are 2 oxygen in reactants and only 1 in products
How can we balance it?
2H2 + O2 → 2H2O

46. Balance the following equations by using
coefficients (“numbers in front of substances”):
a. Ag S  Ag2S
2Ag S  Ag2S
b. Al HCl  H AlCl3
2Al HCl  3H AlCl3
c. KI  K I2
2KI  2K I2
d. K O2  K2O
4K O2  2K2O

47. Agenda for 12/17/15 BELLRINGER: BALANCE THIS PROBLEM
Zn HCl  ZnCl H2
Go over homework
Types and Signs of Chemicals Reactions
Types of Chemical Reactions worksheet
HW – Chemical Reactions worksheet #2
TEST TUESDAY

48. More Balancing Equations
2Na + Cl2 → 2NaCl
2H2 + O2 → 2H2O
CH4 + 2O2 → CO2 + 2H2O
2K + Cl2 → 2KCl
N2 + O2 → 2NO

49. More Balancing Equations
6) Zn + 2HCl → ZnCl2 + H2
7) N H2 → 2NH3
8) Na2SO4 + CaCl2 → CaSO NaCl
9) 2C2H O2 → 4CO H2O
10) 3Mg + N2 → Mg3N2

50. Signs that a chemical reaction has taken place:
1) Change in color

51. 2) Formation of precipitate (formation of a solid when there was none)

52. 3) Release of gas

53. 4) Light is released

54. 5) Energy changes occur
a) Exothermic reaction – gives off energy; gets warmer
Example: burning of wood, explosion of dynamite, Gummy Bear demo, Genie in a Bottle demo
b) Endothermic reaction – absorbs energy; gets cooler
Example: dissolving Epsom salts, reaction of Ba(OH)2 and NH4Cl

55. Types of Chemical Reactions:
1) Synthesis – 2 or more substances combine to form a single, complex product
Examples:
2Mg + O2 → 2 MgO
4Fe O2 → 2Fe2O3 (rust)

56. 2) Decomposition – a complex substance is broken down into two or more simpler substances (opposite of synthesis)
Example: Splitting water by electricity
2H2O → 2H2 + O2

57. 3) Single Replacement – a free element replaces an element in a compound
Example:
Cu + 2AgNO3 → 2Ag + Cu(NO3)2

58. 4) Double replacement – when two ions switch places in a reaction
Example:
Pb(NO3) KI → PbI KNO3