Download Chapter 1 Notes - Spring Branch ISD

Chapter 1 Notes
Scientific Notation
Convert the following numbers to scientific notation
17600.0
4.76
0.00135
-0.1544
10.2
301.0
-67.30
-0.000130
How to expand scientific notation
1) If the exponent is ___________, move the decimal to the _______ to make the number _________.
2) If the exponent is ___________, move the decimal to the _______ to make the number _________.
Samples:
1) Expand the scientific notation of 8.02 x 10-4 to regular notation.
2) Expand the scientific notation of -9.77 x 105 to regular notation.
Expand the following numbers to regular notation
4.96 x 10-2
7.01 x 100
5.50 x 104
5.61 x 10-3
-9.3 x 10-3
4.92 x 102
-8.37 x 101
-9.23 x 10-1
No Calculator Practice
To multiply numbers in scientific notation without a calculator
1) Multiply the number in front
2) Add the exponents
3) Put the answer in correct scientific notation if necessary
To divide numbers in scientific notation without a calculator
1) Divide the number in front
2) Subtract the exponents
3) Put the answer in correct scientific notation if necessary
(2.3 x 103)(1 x 102) =
(6 x 10-7) / (3 x 10-8)
(4.0 x 105) / (2 x 103) =
(2x103)(4x10-2) / (1x10-4)
27 x (3.4 x 102)
(9x10-7) / (1x105)(3x10-1)
Metric Prefixes
Prefix
Abbreviation
Scientific
notation
equivalent
Prefix
Tera
Centi
Giga
Milli
Mega
Micro
Kilo
Nano
deci
Pico
Abbreviation
Convert 4.98 pL to L
Convert 87 µm to Mm
Convert 56 mm to meters
Convert 0.006 kL to nL
Scientific
notation
equivalent
Uncertainty of Measurement
Accuracy = How close to the ___________________ you are.
Precision = How close your measurements are _______________________________.
Rules for counting significant figures
1. Nonzero integers always count Ex 456
_____ sig figs
2. Leading zeros never count.
Ex 0.0025
_____ sig figs
3. Captive zeros always count.
Ex 1001
_____ sig figs
4. Trailing zeros sometimes count Ex 1020
_____ sig figs
1.00
_____sig figs
5. Exact numbers have an infinite number (because it’s not a measurement) Ex 5 students
How many significant figures
20600
0.00960
10 apples
3.90 x 106
10.90
Rules for math functions
1. Multiplication or division
The answer should have the same number of _______________________ as the _________ precise
measurement.
Ex.
1.34 x 0.04 = ______
2. Addition or subtraction
The answer should have the same number of ______________________ as the _________ precise
measurement.
Ex.
1.34 + 2.3 = _______
Practice: The solution to 9.99/22.41 x (18.465 + 0.464) is _________.
Density
__________
The denser object will be on the _______________.
Practice:
1) Water has a density of 1 g/mL. If you have 150mL, how many mg does it weigh?
2) You have an unknown object that weighs 60g. You place it in a graduated cylinder which has 30mL of
water in it. When you place the object in the water the final volume of the graduated cylinder of reads
75mL. What is the density of this unknown object?
Temperature Conversions
Temperature in Celsius + 273 = Kelvin Temperature
o
C
Kelvin
o
C
56oC
Kelvin
190 K
780 K
-93oC
Classification of Matter
Element – matter composed of __________________________ of atom.
Compound – pure substances that are ______________ joined that consist of ___________________
of atom.
Homogeneous Mixture – a ___________ mixture of two or more substances which are
__________________ bonded.
Heterogeneous Mixture – mixture that is __________________ in composition.
Element
Compound
Copper
Chocolate Chip Cookie
Mercury
Coke
Sodium chloride
Vegetable soup
Rocky Road Ice Cream
Homogeneous Mixture
Heterogeneous Mixture
Oil and water
Iron
Rust
Tap water
Milk
Ways to separate a mixture
1) ______________= separating a solution due to differences in boiling points
2) ______________ = Separates a solid and liquid mixture, or a mixture of two liquids at are immiscible.
3) ______________ = Used to separate mixtures by their polarity
Chapter 2 Notes
Atomic Structure
Protons are located in the ___________ and have a __________ charge.
Electrons are located in the ______________ and have a _____________ charge.
Neutrons are located in the ____________ and have a ____________ charge.
Atomic Number = ______________________________________
Atomic Mass = ________________________________________
Chemical Symbol
23
11
23 11 Write the chemical
symbol for
Magnesium-23
Isotope = Atoms that contain the same number of __________________ but a different number of
________________.
12
_______
_______
Ex: C-12 and C-14
14
_______
_____
Atoms are neutral due to equal _______________ and _____________________.
Ions have charges.
Cations = ____________________
Anions = __________________
Practice:
1) How many electrons does Ca2+ have?
__________
2) How many protons does iron have?
__________
3) How many electrons does F-1 have?
__________
4) How many total protons are found in two molecules of C20H30O?
__________
Bonding and Naming
______________ = One loses electrons and one gains electrons. Occurs between metals and
nonmetals.
______________ = Sharing of electrons. Occurs between two non-metals.
Use the naming tree to review your naming rules.
How to name acids if you’re given their formulas:
In summary for acids:
Examples:
1) HNO3
__________________________
____ide hydro ____ic acid
____ate _____ ic acid
2) HF
__________________________
3) H2C2O4
__________________________
4) Hydrosulfuric acid
______________
5) Chromic acid
______________
____ite _____ ous acid
Prefixes:
1- mono
2- di
3- tri
4- tetra
5- penta
6- hexa
7- hepta
8- octa
9- nona
10- deca
Remember that polyatomic ion
names are not changed in any way.
(do not change ending to –ide)
Practice Naming:
1) NaF ____________________
4) Li3PO4
____________________
____________________
5) MgO
____________________
3) H2SO4 ____________________
6) CuO
____________________
2) NO2
Practice Formula Writing:
A. Iron (III) chloride
_____________
D. Hydrofluoric acid
_____________
B. Calcium sulfite
_____________
E. Zinc chloride
_____________
C. Sulfur hexafluoride
_____________
Chapter 3 Notes
Mass spectrometer = tells you which _________________________ are in your sample.
Average Atomic Mass = is found on the periodic table, it is the average of all of an element’s ______________.
= (mass isotope 1) (% abundance isotope 1) + (mass isotope 2) (% abundance 2) + ……
Practice:
1) Assume that element Uus is synthesized and that it has the following stable isotopes:
284 Uus (283.4 amu) 34.60%
285 Uus (284.7 amu) 21.20%
288 Uus (287.8 amu) 44.20%
What would the average atomic mass be?
2) What is the average atomic mass of an element with two isotopes, one isotope has a mass of 64 and is
20% abundant. The other isotope has a mass of 55 and is 80% abundant.
Moles
1 mole = molar mass
1 mole = 6.022 x 1023 atoms/ molecules
1 mole = 22.4 liters
Practice:
1) Calculate the molar mass of calcium chloride
2) How many moles are in 496g of CH4?
3) What is the molar mass of iron (III) chloride
4) How many liters are in 15 moles of oxygen gas
5) How many molecules are in 54.9g of NaCl?
6) How many liters would 100g of CaCl2 take up?
% composition
Part/whole x100
Practice:
1) Calculate %H in water
2) Calculate %N in ammonia
3) What is the % composition of N in nitrogen dioxide?
Empirical Formula = The simplest whole number ratio of atoms in a compound.
Steps to Determine the Empirical Formula
1) Convert % to decimal form for each element
2) Divide decimal form by molar mass
3) Divide each answer by the smallest #
4) Multiply all by a # (if necessary) to get all elements to a whole.
Example:
1) Determine the empirical formula of a compound that is 71.65% Cl, 24.27% C, and 4.07% H.
2) NutraSweet is 57.14% C, 6.16% H, 9.52% N, and 27.18% O. Calculate the empirical formula of
Nutrasweet.
3) Suppose a substance has been prepared that is composed of carbon, hydrogen, and nitrogen. When
0.1156g of this compound is reacted with oxygen, 0.1638g of CO2 and 0.1676g of water is collected.
Assume all Carbon is converted to CO2. I started with how much carbon? What is the empirical formula
for the compound?
4) What is the empirical formula of the compound N6O10?
Molecular Formula = the exact formula of a molecule
_______________ of the empirical formula
Divide the empirical formula’s molar mass by the molar mass of the compound.
Practice:
1) Calculate the molecular formula of a substance if the empirical formula is ClCH2 and has a molar mass
of 98.96 g/mol.
2) A white powder is analyzed and found to contain 43.64% P and 56.36% O by mass. The compound has
a molar mass of 283.88 g/mol. What are the compound’s empirical and molecular formulas?
3) Determine the molecular formula of a compound that contains 26.7% P, 12.1%N, 61.2% Cl and a
molecular weight of 580.
Balancing Equations = you can only add coefficients to make the atoms on the reactants equal the
atoms on the products side.
_____Ca(OH)2 + _____H3PO4 _____H2O + _____Ca3(PO4)2
_____Al(OH)3 + _____HCl _____AlCl3 + _____H2O
_____Au2S3 + _____H2 _____Au + _____H2S
Converting one substance to another = Stoichiometry
1) Balance equation
2) Convert known mass to moles
3) Use appropriate mole ratios to convert moles of given substance to moles of other substance.
4) Convert back to grams if required
_____CH4 + _____O2 _____CO2 + _____H2O
If you want to produce 30g of water how much methane (CH4) should you start with?
_____H2 + _____O2 _____H2O
If you have 5g of oxygen, how much water can you produce if you have excess hydrogen?
_____C7H6O3 + _____C4H6O3 _____C9H8O4 + _____HC2H3O2
What mass of C4H6O3 is needed to completely consume 1.0 x 102g C7H6O3?
What is the maximum amount of C9H8O4 that could be produced if you started with 1.0 x 102g C7H6O3?
Limiting reactant (Reagent) = the reactant that is consumed _______ so it limits the amount of
products that can be made.
Solve the problem twice, once with each reactant, the smaller answer will be the correct
answer.
N2 + 3H2 2NH3
Calculate how much NH3 you would produce if you started with 15g of nitrogen and 25g of hydrogen.
_____H2 + _____O2 _____H2O
Determine how much water you would produce if you started with 6g of hydrogen and oxygen gas.
Hg + Br2 HgBr2
What mass of HgBr2 can be produced from the reaction of 10g Hg and 9g Br2? What mass of which
reactant is left over?
COCl2 + 2NaOH 2NaCl + H2O + CO2
Identify the limiting reagent, and the number of grams left over in excess, respectively, for the above
equation, if 1.9g of each phosgene (COCl2) and sodium hydroxide are combined.
% _______________ !100
Actual yield = _______________________
Theoretical yield = ___________________
2H2 + CO CH3OH
Suppose 68.5 kg CO is reacted with 8.6kg H2. Calculate the theoretical yield of methanol (CH3OH) if
3.57 x 104g CH3OH is actually produced, what is the % yield.
Consider the following unbalanced reaction:
_____P4 + ____F2 _____PF3
How many grams of fluorine are needed to produce 120g of PF3 if the reaction has a 78.1% yield?