Review of Chemistry1.doc

Review of Chemistry
Derived Quantities
Definitions: A DERIVED QUANTITY is a number made by combining two or more other values.
A DERIVED UNIT is a unit which is made by combining two or more other units.
II.4 Density
Definition: Mass= the quantity of matter on an object
Density= the mass contained in a given volume of a substance
II.5 Significant figures and experimental uncertainty
SIGNIFICANT FIGURES
A. A significant figure is a measured or meaningful digit.
B. An accurate measurement is a measurement that is close to the correct or accepted value; a precise measurement is a reproducible measure.
C. The number if significant figures is equal to all the certain digits plus the first uncertain digit.
D. “Defined” numbers and “counting” numbers are assumed to be perfected so that they are “exempt” from the rules applying to significant figures.
BE VERU CAREFU; WHEN A VLAUE APPEARS TO COINCIDE EXACTLY WITH A MARKING ON A MEASURING DEVICE. The following procedure should help when such a situation occurs.
THE PROCEDURE FOR CORRECTLY READING MEASURING SCALES WHEN A POINTER IS EXACTLY ON A NUMBERED DIVISION
l Determine the value that the measurement seems to have.
l Pretend you have a value in between two of the unnumbered subdivisions on your measuring device.
l Determine how many decimal places you could read off the measuring device at the “in-between value”.
l Add a sufficient number of zeroes to the actual reading to give you the correct number of decimal places for your reading.
EXPERIMENTAL UNCERTAINTY
Definition: The experimental uncertainty is the estimated amount by which a measurement might be in error.
E. When adding an uncertainty to a measurement, the uncertainty goes after the measured value but before the unit.
NOTE: If the uncertain digit is in the first decimal place, the uncertainty will be in the first decimal place also.
INTERPRETATION OF UNCERTAINTIES
IMPORTANT: The place value (tens, units, first decimal, etc.) of the experimental uncertainty and the first uncertain digit of a measurement must agree with each other.
F. NORMALLY USE UNCERTAINTIES TO THE NEAREST 0.1 OF THE SMALLEST UNNUMBERED SUBDIVISION. If you can only estimate a value to the nearest 0.2 or even 0.5 of the smallest unnumbered subdivision, feel free to do so, but be prepared justify your decision.
G. Leading zeroes are not significant.
H. Trailing zeroes are all assumed to be significant and must be justified by the precision of the measuring equipment.
There are two ways to count the number of significant figures.
EXPRESS THE NUMBER IN SCIENTIFIC NOTATION AND THEN COUNT ALL THE DIGITS.
Or even simpler:
Starting from the left side of the number, ignore all “leading zeroes” and only start counting at the first NON-ZERO” digit. Once you start counting, continue until you run out of digits.
I. After MULTIPLYING or DIVIDING numbers, round off the answer to the LEASR NUMBER OF SIGNIFICANT FIGURES contained in the calculation.
J. after ADDING or SUBTRACTING numbers, round off the answer to the LEAST NUMBER OF DECIMAL PLACES contained in the calculation.
UNIT III: THE PHYSICAL PROPERTIES ANG PHYSICAL CHANGES OF SUBSTANCES
III.1. some basic definition in science
The following are general characteristics of HYPOTHESES
1. Hypotheses are normally single assumptions.
2. Hypotheses are narrow in their scope of explanation.
3. When originally proposed, hypotheses are tentative but may become generally accepted after more complete testing.
The following are general characteristic of THEORIES
1. Theories are composed if one or more underlying hypotheses.
2. Theories are broad in scope and may have subtle implication which are not foreseen when they are proposed because they provide explanation for entire “fields” of related behavior.
3. Theories are sometimes called models because they often provide a concrete way to examine, predict and test the workings of nature.
4. A theory can’t be “proven” but at some point it may have such a tremendous record of explanation and prediction that we [place a high probability in its correctness as a model capable of describing reality.
5. Theories must be “falsifiable”, they must make testable prediction about the behavior of the system under NEW conditions.
The following are general characteristics of LAWS
1. Laws summarize the results of many experiments or observations and state what will happen when a specific situation occurs.
2. Laws do NOT try to explain WHY something occurs.
3. Laws are NOT “proven theories”, as sometimes is erroneously stated. Laws are often stated before any theory exists as to WHY the law is true.
III.2 THE PHYSICAL PROPERTIES OF MATTER
Definition: MATTER= anything that has mass and occupies space.
(a) Solids are rigid, do not readily change their shape, and experience very small changes in volume when heated or subjected to pressure.
(b) Liquids conform to the shapes of their containers and experience only slight changes in volume when heated or subjected to pressure.
(c) Gases conform to the shapes of their containers and experience drastic changes in volume when heated or subjected to pressure.
An investigation of: vapor pressure, evaporation rate, viscosity, and gas compressibility.
III.3 The classification of matter
1. A HOMOGENEOUS substance is substance consisting of only one phase.
Examples: air, water, salt water, a piece of iron.
2. A HETEROGENENOUS substance is a substance consisting of more than one phase.
Examples: a human being, a pencil. gravel
3. A PURE SUBSTANCE is a substance that is homogeneous and has an unchangeable composition.
Examples: sugar, water, copper, iron
4. A MIXTURE is a system made up of two or more substance, such that the relative amounts of each substance can be VARIED.
Examples: salt dissolved in water, alcohol dissolved in water
“Mixture” is a general term which includes both heterogeneous mixtures (better known as “mechanical mixtures”) and homogeneous mixtures (better known as “solution”).
5. A MECHANICAL MIXTURE is a heterogeneous mixture of two or more substances.
Examples: gravel, sand and iron filings, a pencil
Note: ALL HETEROGENEOUS substances are MECHANICAL MIXTURES, and vice ideas.
6. A SOLUTION is a homogeneous mixture of two or more substances.
There are several different types of solutions.
Types of solution
Example
Gas-in-gas solution
Air (oxygen, nitrogen, etc.)
Gas-in-liquid solution
Soda pop
Liquid-in-liquid solution
Water and alcohol
Solid-in-liquid solution
Salt water
Solid-in-liquid solution
Alloys (metals melted together)
7. As defined previously, an ELEMENT is a substance which cannot be separated into simpler substances as a result of any chemical process. In other words, an element is a pure substance in which all the atoms are of the same types.
8. A COMPOUND is a pure substance made of two or more types of atoms. Only one type of molecule is present in a compound.
The difference in physical properties between different classifications of matter.
III.4 The physical separation of substances
A. Hand Separation
B. Filtration
C. Evaporation
D. Distillation
E. Solvent Extraction
F. Recrystallization
G. Gravity Separation
H. Paper, column, and thin chromatography
A SYNOPSIS OF SEPARATION METHODS
A. Mechanical Mixtures
MIXTURE
METHOD
WHEN TO USE METHOD
SOLID in SOLID
Hand separation
Large chunks present among other solids
Gravity
The density of the desired solids is much different from the density of the other solids
Solvent extraction
One solid preferentially dissolves in a particular solvent
Chromatography
The solids are colored, present in small amounts and are soluble in some solvent or mixture of solvent
LIQUID in LIQUID
Hand separation
A few large pieces of solid are present in the liquid
Gravity separation
Solid particles are present in a small amount of liquid
Filtration
Solid particles are present in a large amount of liquid
B. Solutions
MIXTURE
METHOD
WHEN TO USE METHOD
SOLID in SOLID
Evaporation
The solid is wanted and the liquid is not
Distillation
The liquid is wanted; the solid may or may not be wanted
Solvent extraction
An immiscible added solvent preferentially dissolves at least one but not all of the solids present
Recrystallization
One dissolved solid is much less soluble than the others present (if any); the liquid is not wanted
Chromatography
Small amounts or more than one colored solid are present; the liquid present is not wanted
LIQUID in LIQUID
Distillation
Two or more liquids are present and have different boiling temperatures
Solvent extraction
An immiscible added solvent preferentially dissolves at least one but not all of the liquids present
III.5 Phase changes
IMPORTANT: Chemical changes are frequently accompanied by physical changes. For example, hydrogen gas and oxygen gas react to form liquid water. Therefore, a reference to a chemical change implies that the primary change is chemical, but a physical change may occur as a result.
l On the sloping portions of the graph- all the heat is used to warm the substance so the temperature rises.
l On the level portions of the graph- the substance contains so much heat energy that it cannot absorb more heat and stay in the same phase. The added heat is used, for example, to break up the solid and allow a liquid to form. All the heat is used to change phase so the temperature doesn’t change and the graph levels off.
Special note: When 1g of candle wax is burned about 10000 joules of heat are produced; when 1g of liquid wax solidifies about 40joules of heat are produced. This is typical of the differences between the energy involved in physical changes (for example, the solidification of wax) and chemical changes (for example, the burning of wax). In general: the amount of heat involved in a physical change is much less than the amount involved in a chemical change.
III.6 The role of kinetic energy in physical changes
(a) Rotational energy- causes a molecule to rotate around one of its axes; bond lengths and bond angles don’t change
(b) Vibration energy- changes the bond lengths and/or angles between atoms in a molecule
(c) Translational energy- causes the molecule to travel in a straight line from place to place, but has no effect on bong lengths and angles.
PRACTICAL APPLICATIONS OF KNIETIC ENERGY CHANGES
1. MICROWAVE OVENS supply energy which causes the water molecules in food and liquids to vibrate.
2. IN INFRARED (IR) SPECTROSCOPY, molecules absorb infrared (heat) energy. (“Infrared” is pronounced “infra-red”.) IR energy is just energy with less energy than red light; if white light is passed through a prism, the light energies are spread out as shown below.
Every different molecule has its own “fingerprint”; by comparing the spectrum of an “unknown” molecule with the spectra of known molecules the unknown can be identified.
UNIT IV: INORGANIC NIMENCLATURE
IV.1 the chemical elements
Some comments about the elements may help you to learn their names and symbols.
a) The first letter in the symbol is ALWAYS in upper case (capitals). Second letters, if present, are ALWAYS in lower case.
Example: Pt, Te, Be, Cl
b) Many elements just use the first two letters of the element’s name as the element’s symbol.
Example: Al, Bi, Li
c) When the first two letters have already been used with some other element, the first and the third letters are used.
Example: Ar=argon; As=arsenic (can’t use Ar); At=astatine (can’t use As)
d) Elements which were known in ancient times or which have name taken from substances known in ancient times have symbols derived from their Latin names.
Element
Latin name
Symbol
Antimony
Stibium
Sb
Copper
Cuprum
Cu
Gold
Aurum
Au
Iron
Ferrum
Fe
Lead
Plumbum
Pb
Mercury
Hydragyrum
Hg
Potassium
Kalium
K
Silver
Argentium
Ag
Sodium
Natrium
Na
Tin
Stannum
Sn
e) A few elements have single letter symbols.
Example: B, C, F, H, I, K, N, O, P, S, U, V, W
f) Some elements and their elements frequently cause problems with students. The following might be of some help things straight.
IV.2 naming inorganic compounds
IMPORTANT: Metals form POSITIVE ions.
Nonmetals form NEGATIVE ions (Hydrogen is generally an exception).
Naming monatomic ions
A. naming monatomic metal ions: use the name of the metal and add the word “ion”
Example: Sodium metal (Na) form the sodium ion (Na+)
Aluminum metal (Al) forms the aluminum ion (Al3+)
The stocking System of naming metal ion: If a metal ion has more than one possible charge, the charge is indicated by a Roman numeral, in parentheses, immediately following the name.
Example: Fe3+=iron (III) ion, Fe2+=iron (II) ion, U6+=uranium (VI) ion, U3+= uranium (III) ion
B. naming monatomic non-metal ions: take off the original ending of the element’s name and put on an “ide” ending.
Element name
Element symbol
Ion name
Ion symbol
Fluorine
F
Fluoride
F-
Chlorine
Cl
Chloride
Cl-
Bromine
Br
Bromide
Br-
Iodine
I
Iodide
I-
Oxygen
O
oxide
O2-
Sulphur
S
Sulphide
S2-
Nitrogen
N
Nitride
N3-
Phosphorus
P
Phosphide
P3-
Naming polyatomic ions
Constructing the formula of an ionic compound, given the name of the compound
Constructing the name of an ionic compound, given the formula of the compound
Naming hydrated
Naming compounds by using the prefix-naming system
Some common acids
IV.3 Extension: the colors of the some common aqueous ions
Ion
Color
Ion
Color
Ion
Color
Fe2+
Pale green
MnO4-
Purple
Cr2O72-
Bright orange
Fe3+
Dull yellow
Ni2+
Bright green
Cu2+
Blue(**)
Co2+
Pink-red(**)
CrO42-
Bright yellow
Mn2+
Very pale pink(**)