1. Prof.Dr.Hassan A. Mohammed
Preparatory Program in Basic Science(PPBS001)
L3 ,L4) )Chemistry Course
Instructor:
Prof.Dr.Hassan A. Mohammed

2. Lecture.3 Physical and chemical properties
Properties that describe the look or feel of a substance, such as color, hardness, density, texture, and phase, are called physical properties. Every substance has its own set of characteristic physical properties that we can use to identify
that substance.
Gold
Opacity : opaque
Color : yellowish Phase at 25'C: solid
Density :19.3g/mL
Diamond
Opacity: transparent Color: colorless Phase at 25"C: solid Density: 35 g/mL
Water
Opacity: transparent Color: colorless
Phase at 25"C: liquid Density: 1.0 g/mL

4. The physical properties of a substance can change when conditions change, but that does not mean that a different substance is created. Cooling liquid water to below 0C causes the water to transform to solid ice, but the sub­stance is still water, no matter what the phase. The only difference is the relative orientation of the H20 molecules to one another.
In the liquid phase, the water molecules tumble around one another, whereas in the ice phase, they vibrate about fixed positions. The freezing of water is an example of what chemists call a physical change. During a physical change, a substance changes its phase or some other physical property, but not its chemical composition

5. Gas Liquid Solid Disorder Some space Particles closer together Order
Particles fixed
in position
Total disorder
Lots of empty space
Gas Liquid Solid

6. Chemical Properties
Chemical properties characterize the ability of a substance to transform to a different substance ,normally through a chemical reaction
Chemical Change
1. A chemical reaction transformation, or change
involves a change in the way the atoms in the molecules are chemically bonded to one another.
2. A chemical bond is a force of attraction between two atoms that holds them together.
3. When molecules undergo chemical change, we say that they reacted.

7. Chemical change :
Any change in a substance that involves a rearrangement of the way atoms are bonded is called a chemical change. Examples: Thus the transformation of methane to carbon dioxide and water is a chemical change, CH O → CO2 + 2H2O

8. Baking soda has the chemical property of reacting with vinegar to produce carbon dioxide and water. This is accompanied by the absorption of a small amount of heat energy.
Na2CO CH3COOH CH3COONa + CO2+ H2O
(c) Copper has the chemical property of reacting with carbon dioxide and water to form substance known as patina greenish-blue solid

9. (d) When an electric current is passed through water, its energy causes a reaction in which the water molecules split into atoms, which then recombine into hydrogen and oxygen molecules that bubble up to the surface

10. In the language of chemistry, materials undergoing a chemical change are said to be reacting. Methane reacts with oxygen to form carbon dioxide and water.
Water reacts when exposed to electricity to form hydrogen gas and oxygen gas.
Thus, the term chemical change means the same thing as chemical reaction. During a chemical reaction, new materials are formed by a change in the way atoms are bonded together.

11. physical and chemical changes.
There are two powerful guidelines that can assist you in assessing physical and chemical changes.
First in a physical change, a change in appearance is the result of a new set of conditions imposed on the same material. Restoring the original conditions restores the original appearance: frozen water melts upon warming.
Second, in a chemical change, a change in appearance is the result of the formation of a new material that has its own unique set of physical properties.
The more evidence you have suggesting that a different material has been formed, the greater the likelihood that the change is a chemical change. Iron is a material that can be used to build cars. This suggest that the rusting is a chemical change

12. Examples For Chemical And Physical Changes:
1-Potassium chromate, at room temperature is a bright canary yellow. At higher temperature, it is a deep reddish orange. Upon cooling the canary yellow color returns, suggesting that the change is physical.
With a Physical change, reverting to the original
conditions does restore the original appearance.
2-Ammonium dichromate, is an orange material that when heated explodes into ammonia, water vapor, and green chromium (lll) oxide. A return to the original temperature does not restore the orange color. This is a chemical change.

13. Lecture. 4 Elements to compounds and naming compounds 2
Lecture.4 Elements to compounds and naming compounds 2.3 Forming compounds
A molecule is a particle made from two or more atoms. These atoms can be identical (as for H2, N2, O2, etc.), or different (H2O –NaCl..).
A compound is a collection of identical molecules, each consisting of more than one element. A molecule, therefore, is the smallest amount of a compound.

14. Examples:
When atoms of different elements bond to one another, they make a compound. Sodium Na atoms and chlorine Cl atoms, for example, bond to make the compound sodium chloride NaCl, , commonly known as table salt.
Nitrogen N atoms and hydrogen H atoms join to make the compound ammonia NH3 , which is a common household cleaner.

15. The sodium chloride, NaCl, shown in Figure is very different from the elemental sodium and the elemental chlorine used in its formation. Elemental sodium, Na, consists of nothing but sodium atoms, which form a soft, silvery metal that can be cut easily with a knife.
(a)Sodium metal (immersed in oil to prevent reaction with oxygen and moisture in the air); (b) chlorine gas; (c) the reaction between sodium and chlorine; (d) sodium chloride (common table salt)

16. Naming Compounds
A system for naming the countless number of possible compounds has been developed by the International Union for Pure and Applied Chemistry (IUPAC). This system is designed so that a compound's name reflects the elements it contains and how those elements are joined together.

17. NaCl Sodium chloride HCI Hydrogen chloride
Guideline 1.
The name of the element farther to the left in the periodic table is followed by the name of the element farther to the right, with the suffix: " ide" added to the name of the latter:
NaCl Sodium chloride HCI Hydrogen chloride
Li2O Lithium oxide MgO Magnesium oxide
Calcium fluoride CaF2 Sr3P Strontium phosphide

18. Guideline 2
When two or more compounds have different numbers of the same elements, prefixes are added to remove the ambiguity. The first four pre­fixes are "mono-" (one), "di-" (two), "tri-"-(three), and "tetra-" (four). The, prefix mono-, however, is commonly omitted from the beginning of the first word of the name
Carbon and oxygen
CO Carbon monoxide
CO Carbon dioxide
Nitrogen and oxygen
Nitrogen dioxide NO2
Dinitrogen tetroxide N2O4

19. Guideline 3. Many compounds are not usually referred to by their systematic names. Instead, they are assigned common names that are more convenient pr have been used traditionally for many years. Some common names are water for H20, ammonia for NH3 and methane for CH4.

20. شكراً لحسن أستماعكم

22. Prof.Dr.Hassan A. Mohammed
Preparatory Program in Basic Science(PPBS001)
L5 ,L6) )Chemistry Course
Instructor:
Prof.Dr.Hassan A. Mohammed

23. Mixture, classification of matter and solutions
Lecture.5
Mixture, classification of matter and solutions
2.4 Most Materials Are Mixtures
A mixture is a combination of two or more substances in which each substance retains its own properties.
Most materials we encounter are mixtures:" mixtures of elements, mixtures of compounds, or mixtures of elements and compounds. Examples :
(a) Stainless steel, for example', is a mixture of the elements iron. Chromium, nickel and carbon.
(c) Tap water is a location-dependent mixture. It contains mostly water, but also many other compounds such as calcium, magnesium, fluorine, iron, potassium, and traces of lead, mercury, and other materials. .

24. (c) Our atmosphere, , is a mixture of the elements nitrogen; oxygen, and argon, plus small amounts of such compounds as carbon dioxide and water vapor
(d) In the figure, the sugar molecules in the
teaspoon of sugar are identical to those already in the tea. The latter are mixed with other substances, mostly water

25. We may separate a mixture’s components by taking
advantage of the difference in their physical properties(such as melting or evaporation, boiling points, molecule size, etc.)
Examples:
(a) We can separate H2O from NaCl in seawater by distillation ,which is boiling and evaporating the water.
(b) A solid-liquid mixture (such as coffee) can be separated by filtration, which is to use a filter that passes the liquid and catches the solid particles.

27. The chemist's classification of Matter

28. :Solutions & Lecture 6 : Mixtures
Mixtures can be classified as heterogeneous or homogeneous.
A homogeneous mixture is one in which the substances are evenly distributed .Tap water is a homogeneous mixture, for example.
A heterogeneous mixture contains substances that are not evenly distributed, so different regions of the mixture have different properties.
A bowl of cereal is an obvious heterogeneous mixture.
A suspension is a homogeneous mixture in which the different components are in different phases (such as solids in liquids or liquids in gases). Examples: (a) Milk is a suspension of proteins and fats (solid) finely dispersed in water (liquid). (b) Clouds are a suspension of water droplets (liquid) suspended in air (gas).

29. Solutions
A solution is a homogeneous mixture of molecules (or ions). At the submicroscopic level, a solution is a finely mixed homogeneous mixture of atoms and molecules.
Examples:
(a) Sugar in water is a solution in the liquid phase.
(b) Metal alloys are solid solutions, mixtures of different metallic elements. The alloy brass is a solid solution of copper and zinc, for instance, and the alloy stainless steel is a solid solution of iron, chromium, nickel, and carbon. Rose gold and white gold are other
examples of solutions in the solid state.

30. c) An example of a gaseous solution is the air we inhale
c) An example of a gaseous solution is the air we inhale. By volume ,this solution is 78% nitrogen gas, 21% oxygen gas, and 1% other gaseous materials, including water vapor and carbon dioxide. The air we exhale is a gaseous solution of 75% nitrogen, 14% oxygen, 5% carbon dioxide, and around 6% water vapor.
In describing solutions, the component present in the largest amount is called solvent ,and the other components are called solutes.
The process of a solute mixing in a solvent is called dissolving. In the above example of tea, sugar is dissolved in water, making sugar a solute and water the solvent.

32. Some Common Solutions Chemistry 140 Fall 2002
Liquid solutions are most common.
Solid solutions – alloys
Prentice-Hall © 2002

33. Notes A material’s readiness to dissolve in another is a function of
the electrical attraction between them.
2. There is a limit to how much of a given solute can dissolve in a given solvent.
Example: When you add a little sugar to a glass of water, the sugar rapidly dissolves. As you continue to add sugar, however, there comes a point when it no longer dissolves. Instead, it collects at the bottom of the glass, even after stirring. At this point, the water is saturated with sugar, which means that it cannot accept any more sugar. When this happens, we have what is called a saturated solution, defined as one in which no more solute can dissolve.
A solution that has not reached the limit of solute that will dissolve is called an unsaturated solution.

34. The quantity of solute dissolved in a solution is described in mathematical terms by the solution's concentration which is the
amount of solute dissolved per amount of solution:
amount of solute
Solution’s concentration =
amount of solution
Example:
A sucrose-water solution may have a concentration of 1 gram of sucrose for every liter of solution (1 g/L). A 2-g/L sucrose-water solution is more concentrated, and 0.5-g/L solution is less concentrated, or more dilute (!). We are often more interested in the number of solute particles in a solution rather than the number of grams of solute. For this purpose, we use a unit called the mole. By definition:
1 mole (of any type of particles) = 6.02 × 1023 particles

35. This number is known as Avogadro's
number (NA), after the 19th-Century
Italian scientist who first proposed it.
The molar mass (mass of 1 mole, or
6.02 × molecules) of a particular
type of molecules is calculated as the
sum of the atomic masses (in grams) of
all its constituent atoms.
Example:
Sucrose has the chemical
formula: C12H22O11. Its molar mass is:
Molar mass of sucrose = (12g×12) + (1g×22) + (16g×11) = 342 g
This makes about a cupful of sugar. Thus, we say that grams of
sucrose contain 1 mole of sucrose.

36. A common unit of concentration used by chemists is molarity, which is the solution’s concentration expressed in moles of solute per liter of solution (moles/L):
number of moles of solute
Molarity =
liters of solution
A 1-mole/L concentration is also called1 molar (or 1 M). A 2-M solution has double this concentration.
Worked Exercise:
A saturated aqueous solution( C6H12O11 ) contains 200 g of sucrose and 100 g of water. Which is the solvent: sucrose or water?
Mole mass of water (H2O) = (1g×2) + (16g×1) = 18 g
Molecules of sucrose in solution = (200/342) × × = ×1023

37. Molecules of water in solution = (100/18) × 6.02 × 1023 = 33.4 × 1023
As defined earlier, the solvent is the component present in the largest amount, but what do we mean by amount? If “amount” means mass, then sucrose is the solvent. If “amount” means number of molecules (usually, a chemist’s viewpoint), then water is the solvent.
A unit used for very low (but significant) concentrations is: parts per million (ppm).
A concentration of 1 ppm means that there is 1 particle
of substance for a million total particles. For example, a drop of orange juice in a 40-gallon barrel of water would have a concentration of 1 ppm.

38. Expressing Solution Concentration
Mass Percentage:
Parts per million (ppm):
Parts per billion (ppb):

39. Molality (m – mols/kg):
Mole fraction (X):
Molarity (M – mols/L):
Because volume is temperature dependent, molarity can change with temperature.
Molality (m – mols/kg):
Because both moles and mass do not change with temperature, molality (unlike molarity) is not temperature dependent.

40. شكراً لحسن أستماعكم

41. Prof.Dr.Hassan A. Mohammed
Preparatory Program in Basic Science(PPBS001)
L7 ,L8) )Chemistry Course
Instructor:
Prof.Dr.Hassan A. Mohammed

42. Electron-dot structure and formation of ions
Lecture.7
Electron-dot structure and formation of ions
Lewis structures, also called Lewis-dot diagrams, Electron-dot diagrams or Electron-dot structures, are diagrams that show the bonding between atoms of a molecule, and the lone pairs of electrons that may exist in the molecule.
A Lewis structure can be drawn for any covalently-bonded molecule, as well as coordination compounds.
The Lewis structure was named after Gilbert N. Lewis, who introduced it in his 1916 article “The Atom and the Molecule”. They are similar to electron dot diagrams in that the valence electrons in lone pairs are represented as dots, but they also contain lines to represent shared pairs in a chemical bond (single, double, triple, etc.).

43. Lewis structures show each atom and its position in the structure of the molecule using its chemical symbol. Lines are drawn between atoms that are bonded to one another (pairs of dots can be used instead of lines). Excess electrons that form lone pairs are represented as pairs of dots, and are placed next to the atoms.
A Lewis structure is a type of shorthand notation. Atoms are written using their element symbols. Lines are drawn between atoms to indicate chemical bonds. Single lines are single bonds. Double lines are double bonds. Triple lines are triple bonds. (Sometimes pairs of dots are used instead of lines, but this is uncommon.) Dots are drawn next to atoms to show unbonded electrons. A pair of dots is a pair of excess electrons.

44. Bond and Lone Pairs
Valence electrons are distributed as shared or BOND PAIRS and unshared or LONE PAIRS. • •• H Cl
shared or
bond pair
lone pair (LP)
This is called a LEWIS structure.

45. When compounds are formed they tend to follow the Octet Rule.
Octet Rule: Atoms will share electrons (e-) until it is surrounded by eight valence electrons.
Rules of the (VSEPR) game:
i) O.R. works mostly for second period elements.
Many exceptions especially with 3rd period elements (d-orbitals)
ii) H prefers 2 e- (electron deficient)
iii) :C: N: :O: :F:
4 unpaired 3unpaired 2unpaired unpaired up = unpaired e-
4 bonds 3 bonds 2 bonds 1 bond
O=C=O N=N O = O F - F
iv) H & F are terminal in the structural formula (Never central)

46. Lewis Dot Diagrams of Selected Elements

47. Formation of ammonium cation
Formation of Oxygen molecule