Rules For Significant Figures



Scientific Notation

• If there is no decimal place a decimal at the end of the number

• Move the decimal so that the new number is between 1-10

• Write the new number and place a (x10) after the number

• Place an exponent above the 10, the exponent should be the number of places the decimal was moved.

• If the original number was greater than 1, than the exponent will be positive. If the number is less than one, the exponent will be negative

Example: 45,000 = 4.5 x 104

0.00045 = 4.5 x 104

** Use EE button on calculator for x10

________________________________________________________________________

Rules for Significant Figures

No decimal: Count Right to Left- 24,500 count ( has 3 significant figures

Decimal: Count Left to Right- 0.045 count ( has 2 significant figures

All non-zero digits (1-9) are significant

ZEROS WHICH ARE PLACE HOLDERS ARE NOT SIGNIFICANT

All non-zero digits (1-9) are significant.

Zeros between non-zero digits are significant. Ex: 3004 has 4 significant figs.

Zeros to the left of the first non-zero digit are not significant. Ex: 0.004 has 1 significant fig.

If a number ends in zeros to the right of the decimal point, those zeros are significant. Ex: The number 40.0 has 3 significant figures

Multiplication and division, the result may have no more significant figures than the factor with the fewest number of significant figures.  

Example:  2.52  x  1.003  =  2.52756 ( recorded to 2.53

(3 sig figs) (4 sig figs) = (3 sig. figs) (smallest)

Adding and subtracting: Add or subtract the numbers.  The final answer can contain only as many decimal places as found in the measurement with the fewest number of decimal places.

Example:  134.052 m   +   1.23 m =  135.282 ( rounds to 135.28

(3 decimal places) (2 decimal places) (2 decimal places)

________________________smallest_________________________________________

Density (mass unit/volume unit) = Mass (unit) ÷ Volume (unit) ________________________________________________________________________

Specific heat (c) (units) = calories / (grams x (C) or joules / (grams x (C)

Heat released (cal or J) = mass (g) x Specific heat (choose one unit above) x temp change (ºC) ________________________________________________________________________

Conversions: When converting units make sure the unit you are looking for is on top of the conversion fraction and the unit given is on the bottom of the conversion fraction so the units given can cancel.

1 dozen

Example: 132 pencils = ________ dozen 132 pencils x ---------------- = 11 dozen

12 pencils ________________________________________________________________________

Lewis / Electron-dot diagrams

1. Write symbol of element

2. Place the correct number of dots around the symbol (The number of dots equals the number of valance electrons or the group / roman number the element is in)

3. Spread out the dots before doubling up

●

Aluminum = Group III ( ● Al ●

●●

Nitrogen = Group V ( ● N ●

●

Metric Conversions

|Units (Decreasing |Length |Mass |Volume |

|order) |meters |grams |liters |

|Kilo: |0.001 Km |0.001 Kg |0.001 KL |

|Base unit: |1 m |1 g |1 L |

|Centi: |100 cm |100 cg |100 cL |

|Milli: |1000 mm |1000 mg |1000 mL |

|Micro: |1,000,000 µm |1,000,000 µg |1,000,000 µL |

|Nano: |1,000,000,000 nm |1,000,000,000 ng |1,000,000,000 nL |

| | | |1cm3 = 1 mL |

When choosing the ratios from the table above, one part of the ratio goes on top of the fraction and the other goes on the bottom. The ratio should be set up so that the top of the ratio should have the units that you are looking for and the bottom should be the units that you start with.

Example: 2 mm = _______ cm

Get from chart above

100 cm

2 mm ( ---------------- = 0.2 cm

1000 mm

________________________________________________________________________

Elements

Nucleus is positively charged

Number of neutrons= atomic mass - atomic number

Atomic number= number of protons

Number of protons= number of electrons

Valence electrons= group number

Electron: negative sub particle (outside of the nucleus) CHARGE= -1

Proton: positive sub particle (found in nucleus) CHARGE= +1

Neutron: neutral sub particle (found in nucleus) CHARGE= 0

Periodic table of elements: list of elements in order

______________________________________________________________________________

Percent Error Formula

Percent error = [Measured value – Accepted value] ( Accepted value x 100

Percent Composition

Percents

1. Place each part of the mixture over the total

2. Each fraction is divided

3. Each decimal is multiplied by 100

Percent composition

Formula= Percent composition = (Mass of element x subscript) ( (mass of compound) x 100

• the mass of each element is multiplied by their subscript

• the molar mass of the compound is the sum of the elements

________________________________________________________________________

Writing Formulas

Ionic Compounds: oppositely charged ions. A metal and nonmetal(s)

Binary Ionic Compounds Criss-cross to arrive at subscripts Ex: Sodium nitride = Na+1 and N -3 ( Na3N

Ternary Ionic Compounds

Aluminum hydroxide = Al+3 and (OH-1) ( switch numbers to get Al(OH)3

Binary Molecular (Covalent) Compounds:

Prefixes (see chart below) become subscripts: Ex: Trinitrogen pentabromide ( N3Br5

Naming Compounds

Binary Ionic:

1. One metal and nonmetal 2. If metal is from transitional group must use roman number if not just name metal 3. Nonmetal drops ending and adds “ide”

Example: CaCl2 Calcium chloride

Example: Cr2S3 Chromium (IIII) sulfide (Roman number is found as other subscript)

Ternary Ionic Compounds

1. One metal and more than one nonmetal

2. If metal is from transitional group must use roman number if not just name

3. Nonmetal group gets name from polyatomic ion chart.

Example: Ca(NO3)2 Calcium nitrate

Example: Fe(OH)2 Iron (II) hydroxide (Roman number is found as other subscript)

Binary Molecular (Covalent) Compounds: Both negative charged ions. Both non-metals (Use prefixes as subscripts)

1. Second nonmetal drops ending and adds “ide” Ex: N2O3 Dinitrogen trioxide

|Prefix |Number of atoms |Prefix |Number of atoms |

|Mon |1 |Hexa |6 |

|Di |2 |Hepta |7 |

|Tri |3 |Octa |8 |

|Tetra |4 |Nona |9 |

|Penta |5 |Deca |10 |

Drawing Bohr diagrams

1. The number of protons and neutrons are place in the nucleus.

2. The electrons are placed in energy levels according to the maximum occupancy.

(Fill each shell in numerical order until all electrons are used)

|Energy Level |Maximum number |

| |of electrons |

|1 |2 |

|2 |8 |

|3 |18 |

|4 |32 |

﴿ ﴿ ﴿

Ex: Argon #18 2e- 8e- 8e-

﴿ ﴿ ﴿

Temperature Conversions using formulas

1. To change Celsius to Fahrenheit: Take the given temperature and multiply by 1.80 then

press equal then add 32 then press equal.

2. To change Fahrenheit to Celsius: Take the given subtract 32 then press equal then multiply

the number by 0.56 then press equal.

3. To change Celsius to Kelvin: Add 273 to the given temperature

4. To change Kelvin to Celsius: Subtract 273 from the given temperature.

Types of Equations

Combination: A + B ( AB

Decomposition: AB ( A + B

Single Replacement: A + BC ( AC + B (A is metal)

A + BC ( BA + C (A is a nonmetal)

Double Replacement: AB + CD ( AD + CB

Complete Combustion: AB + O2 ( CO2 + H2O

Incomplete Combustion: AB + O2 ( CO + H2O

________________________________________________________________________

Scientists

JJ Thompson discovered the negative part of the atom

Rutherford discovered the center of the atom

Aristotle believed all matter was made of earth, fire, wind and water

Mendeleev organized the elements

________________________________________________________________________

Atomic Mass

Diatomic Elements: H2, N2, O2,, F2, Cl2, Br2, I2

Double the mass

Example: Bromine: 79.9 x 2= 158.80

Compounds

Multiply the subscript of each element by the element’s atomic mass, and then add.

Example: H2 O

H2= 1.01 x 2 = 2.02

O= 15.99 x 1 = 15.99

Atomic Mass : 18.01

Compounds with Parenthesis

Distribute number outside of parenthesis to numbers inside parenthesis, multiply the subscript of each element by the element’s atomic mass, and then add.

Example: Ca(OH) 2

Ca O2 H2

Ca: 1 x 40.07 = 40.07

O2: 2 x 15.99 = 32.00

H2: 2 x 1.01 = 2.02

74.09

The Mole

Mass to Moles

Take the number of grams given and divide by the mass of the compound from the periodic table

Example: 40.0 grams of H2O=

40g H2O / ( 2xH (1.01) + 1xO (16.0))

40g H2O / (2.02 + 16)

40g H2O/ 18.02= 2.22 moles

Moles to Mass

Take the number of moles given and multiply by the mass of the compound from the periodic table

Example: 4.00 moles of H2O =

4.0 mol H2O x (2 x H (1.01) + 1 x O (16.0))

4.0 x (2.02 + 16) = 72.08 grams

Moles to Particles

Take the number given and multiply by 6.02 x 1023

Example: 4.5 moles of H2O=

4.5 mol H2O x 6.02 x 1023

2.71 x 1024 particles

Particles to Moles

Take the given number and divide by 6.02 x 1023

Example: 4.50 x 1024 particles of H2O=

4.50 x 1024 / 6.02 x 1023= 7.47 moles

-----------------------

18p+

22 n0

................
................

In order to avoid copyright disputes, this page is only a partial summary.

Google Online Preview   Download