Significant Figures(SF) Rules

Precision

        Degree of exactness to which a measurement can be reproduced

        Limited by the finest division on its scale

Accuracy

        Agreement of a particular value with the true value

Significant Figures (SF)

There are 5 rules about significant figures:

1.     Non-zero digits are always significant.

Eg. 1285- 4SF

2.     Zeros between two non- zero numbers are significant.

Eg. 809- 3SF

3.     Zeros at the beginning of a number are never significant.

Eg. 0.02- 1SF

4.     Zeros that fall at the end of a number and after the decimal point are always significant.

Eg.6.100- 4SF

5.     Zeros at the end are Ambiguous, they are not considered significant unless there is a decimal that follows it.

Eg. 600-1SF

Eg. 320.- 3SF

How certain are you?

Absolute Uncertainty
Within every set of data there is an absolute uncertainty. This is the average of the numbers minus the number with the largest difference. Before you can caculate the average though you need to make sure to remove all inprecise data.

For example the numbers 11.9cm, 12cm, 11.8cm, 10.6cm, 11.7cm
You would remove the 10.6cm because it does not match the other data.
Next you would find the average by adding all the numbers that are left then dividing giving you 11.85
Then you would find the number with the largest difference which would be 11.7 and this would give you the answer 11.85 +/- 0.15 cm

Also this can be represented as a percentage by dividing the uncertain value by your average
The lower the percentage the more precise the values were.

Uncertainties in Measuring
When using a measuring tool you can estimate one more decimal place than the actual value shown on the the tool. Like on a ruler that is in mm you can guess one more decimal place further than mm. This value is one tenth of the smallest unit of measuring on a tool. With a 50mL graduated cyclinder your uncertainty would be 5mL.

Separation Techniques

 

Separation Techniques

Basis for separation- different component different properties

Strategy: devise a process that discriminates between components with different properties

Separation

Components in a mixture retain their identities

The more similar the properties are, the more difficult it is to separate them

Basis Techniques

Filtration

Floatation

Crystallization and Extraction

Distillation

Chromatography

Hand separation and Evaporation- Boil away the liquid and the solid remains (solid to solid)

A mechanical mixture of heterogeneous mixture can be separate by using a magnet

Filtration (solids (not dissolved) and liquids)

-Pass a mixture through a porous filter

Crystallization (solid in liquid)

-Precipitation, solids are then separated by filtration or floatation

Saturated solution of a desired solid

Evaporate or cool- solid comes out as pure crystal

Gravity (solids based on density)

-A centrifuge whirls the test tube around at high speeds forcing the denser materials to the bottom. Work best for small volume.

Solvent Extraction

-Mechanical mixture- use liquid to dissolve one solid but not the other so the desired solid is left behind or dissolved

Solution- solvent is insoluble with solvent already present

Distillation (liquid in liquid solution)

-Heating a mixture can cause low boiling components to volatilize (vaporize)

Chromatography

-Flow the mixture over a material that retains some components more than others, so different components flow over the materiel at different speeds

-Sheet

        Paper Chromatography (PC)

-stationary phase is liquid soaked into a sheet or strip of paper

-components appear as separate spots spread out on the paper after drying or “developing”

        Thin layer chromatography (TLC)

Stationary phase is a thin layer of absorbent (Al2O3 or SIO2, usually coating a sheet of plastic or glass)

Quantities and Unit Conversion

 

Quantities and Unit Conversion

Any measurement is always a multiple of some basic unit

Many units describe the same measurement

Quantities

All measure has to have 2 parts: the number and the unit

Numbers and unit combination are called quantities

SI (system International)- A French system dating back to the early 1800s using power of 10s

SI Base Units:

Measure	Unit	Abbacy
Length	Meter	m
Time	Second	s
Mass	Kilogram	kg
Amount of substance	Mole	mol
Luminous intensity	Candela	cd
Temperature	Kalvin	K
Electric Current	Ampere	A

Conversions

For converting meters and cm, we use the equation 1m=100cm since 100cm=1m 1=1m/100cm

Eg 5cm=0.05m

5cm/1 * 1m/100cm=0.005m

The main advantage of including the unit symbols is that you can keep tract of exactly what you are doing with calculation

Text	Symbol	Factor
tera	T	1,000,000,000,000
giga	G	1,000,000,000
mega	M	1,000,000
kilo	k	1,000
hecto	h	100
(none)	(none)	1
centi	c	0.01
milli	m	0.001
micro	μ	0.000001
nano	n	0.000000001

Using Chromatography: Lab

     Chromatography is a way to seperate mixtures to isolate the different components. To seperate the mixtures soluables are added to solvents and then identified depending on they're soluability or being absored into a solid.

     To perform a chromatography experiment you need a carrier phase and a stionary phase. In this case during paper chromatography the sample is spotted on the paper and the sample is carried along by the solvent that acts as the moving carrier. Depending on the solubilities of each component of a mixture, the components will be carried different distances. This allows you to get an Rf  (ratio of fronts) value.


Rf  values can be obtained by diving the distance the solute moved by the solvent front (the distance the solvent travelled). Using Rf  values you can differentiate inbetween two substances and identify them accordingly.


Ex: Solute front (d1 the distance the solute traveled) = 3.1cm
       Solvent front (d2  the distance the solvent traveled) = 8.1cm

        d1/d2=Rf
             3.1/8.1=0.382

       This Rf value corresponds to the color red so it can be red food coloring.

    

Ionic, Covalent, and Acid Compounds

 

How acids are formed

Example:

H(+) + CL(-) = HCL(g) ionic “non acid” hydrogen chloride

HCL(g) + H2O = H3O(+)(aq) + CL(-)(aq)

Hydrochloric acid

1.      Name the positive ion first “hydrogen” (H+)

2.      Name the negative ion second (use the ion names listed in the “table of common ions”)

3.      Remember that the total charge on an ionic compound must be zero. Therefore believe positive and negative charges.

Solutions of hydrogen combined with non metals from groups 16 and 17 and simple acids.

1.      The prefix “hydro” is used as the beginning of the acid name.

2.      The last syllable in the name of the non metal is replaced with the name “ic”

Rules for complex acids

1.      No hydrogen for the ionic non acid name

2.      Ate replace with “ic”

Ite replace with “ous”

“We ate – icy sushi and get appendicite-ous”

Law of Definite Composition (Proust’s Law)

Chemical compound always has the same proportion of elements by mass

Ex: H2O has 2 H and 1 O for a local mass of 18g (H=2g and O=16g) which would apply anywhere in the universe.

Law of Multiple proportion (Dalton’s Law)

Same elements can combine in more than one proportion to form different compounds.

Ex: PbO and PbO2

Matter

Three States of Matter

Solids:

Atoms in solids are tightly packed in an orderly fashion, and vibrate slowly but don’t move from place to place.

- Keeps its own volume and shape.

Liquids:

Atoms are close together and free flowing, sliding past one another.

- Takes the shape of its container.

Gas:

Atoms are spaced far apart with no arrangement and move quickly.

- Takes the shape and volume of its volume.

MIXTURES

-          more than one set of properties

-          physically combined

-          more than one kind of substance

Homogeneous (solutions and colloids)

-          uniform throughout

-          appears to have only one component

Heterogeneous (suspensions and mechanical mixtures)

-          non uniform

-          appears to have more than one component

PURE SUBSTANCE

-          one set of properties

-          one kind of particle

Element (metals, metalloids and non-metals)

-          simplest form of matter

-          can not be decomposed

-          made up of atoms

Compound

-          made up of elements

-          smallest particle is called a molecule

-          ionic (acid, base, salt)

-          covalent (organic compounds)

Chemical and Physical Changes

Chemical & Physical Changes

            It is important to understand the difference between chemical and physical changes. The two types are based on studying chemical reactions and states of matter. 

·         Chemical Changes

               Chemical changes take place on the molecular level. A chemical change produces a new substance. Examples of chemical changes include combustion (burning), cooking an egg, rusting of an iron pan, and mixing hydrochloric acid and sodium hydroxide to make salt and water.

·         Examples of Chemical Changes

1)      burning wood

2)      dissolving salt in water

3)      mixing acid and base

4)      digesting food

·         Physical Changes

              Physical changes are concerned with energy and states of matter. A physical change does not produce a new substance. Changes in state or phase (melting, freezing, vaporization, condensation, sublimation) are physical changes. Examples of physical changes include crushing a can, melting an ice cube, and breaking a bottle.

·         Examples of Physical Changes

1)      crumpling a sheet of paper

2)      melting an ice cube

3)      casting silver in a mold

4)      breaking a bottle

·         How to Tell Chemical & Physical Changes Apart

              A chemical change makes a substance that wasn't there before. There may be clues that a chemical reaction took place, such as light, heat, color change, gas production, odor, or sound. The starting and ending materials of a physical change are the same, even though they may look different.