Unit 1 review

Chapter 1, chapter 2.1-2.2, 2.4

The Organization of Matter Matter

Mixtures:a) Homogeneous (Solutions)b) Heterogeneous

Pure Substances

Compounds Elements

Atoms

Nucleus Electrons

Protons Neutrons

Quarks Quarks

Pure Substances• Each have a fixed composition

• Each a unique set of properties– Elements –

• a type of matter that cannot be broken down into two or more pure substances

– Compounds –• A pure substance that contains more that one element.

Mercury, Hg

Chlorine gas, Cl2

Sodium metal, Na

Sodium Chloride, NaCl

Silver metal, Ag

Water, H2O

Phase Differences

Solid – definite volume and shape; particles packed in fixed positions.Liquid – definite volume but indefinite shape; particles close together but not in fixed positionsGas – neither definite volume nor definite shape; particles are at great distances from one anotherPlasma – high temperature, ionized phase of matter as found on the sun.

Properties of Substances• Chemical properties– Observed when substances take part in a chemical

reaction –• A change that converts it to a new substance.

• Physical properties– Observed without changing the chemical identity of a

substance.• Melting point• Boiling point• Color• Texture• Density• others

Properties of Gold are:

Physical: Melting point of 1063 oC (intensive)

Color gold (intensive)

Amount in weight (extensive)

Chemical: Gold can be stored in air without reacting chemically with oxygen

Separation of a CompoundThe Electrolysis of water

Water → Hydrogen + Oxygen2 H2O → 2 H2 + O2

Reactant → Products

Compounds must be separated by chemical means.

With the application of electricity, water can be separated into its elements

Mixtures• Contain two or more substances combined in such a

way that each substance retains its chemical identity.– Homogeneous

• Uniform mixture in which the composition is the same throughout.

– Heterogeneous• Nonuniform, different in composition throughout.

Granite

HeterogeneousBrass

Homogeneous

Copper Sulfate, heterogeneous

Separation of a Mixture

The constituents of the mixture retain their identity and may be separated by physical means.

Separation of a Mixture

The components of dyes such as ink may be separated by paper chromatography.

Separation of a Mixture by Distillation

Measurements• Scientific measurements are expressed in the

metric system.

• You will need to review this system if you are unfamiliar with it. Pages 14 in you text.

Fundamental SI units(systeme international-units agreed upon by the science community)

• Mass is measured in grams

• Length is measured in meters

• Time is measure in seconds

• Temperature is measured in Kelvin

• Electric currents is measure in amperes

• Amount of substance is measured in moles

All other units are derived from these basic units

TemperatureKelvin = celsius+ 273

Derived units

• Volume measured in cm3

– Derived from finding the volume of a cube

– Length (cm) x width (cm) x height (cm) = cm3

– Common unit (but not the SI) is liters

– The AP test will use liters or milliliters

Density

the ratio of mass to volume.

D = m/V

SI unit: g/cm3

common in g/mL

AP test uses: g mL-1

Significant Figures• Every measurement carries with it a degree

of uncertainty.

• This depends upon what instrument is being used to measure.

• We do not use the + in significant figures however it is understood that there is an uncertainty of at least one unit in the last digit place.

Rules of significant figures• If it is a number 1-9 it is significant.

• If it is a zero between two numbers, it is significant.

• If it is a zero that tells how well something was measured, it is significant. (zeros at the end with a decimal in the number)

• If it is a zero that just tells how big or how small a number is, it is NOT significant. (zeros at the beginning and at the end with no decimal in the number)

Examples

How many significant figures are in each ?•5.00•0.0090•2010•400.0•0.0609•1.50 x 103

•Answers•3•2•3•4•3•3

• When multiplying or dividing significant figures:– Count how many significant figures in each of the numbers

being used. Then use the smallest amount of significant figures of all the numbers when reporting the answer.

• When adding or subtracting significant figures:– Line up the decimals, the last place that is significant in

both numbers is where you draw a line. Add or subtract the numbers, then look at the number just past the line if it is bigger or equal to 5 round up if not just drop after the line.

Examples

Calculate and record the correct number of significant figures.1. 5.60 x 0.0232. 6.03 x 10 23/1.0 x1012

3. 5.1 + 54.674. 455 -22.0

Answers1. 0.132. 6.0 x 1011

3. 59.84. 433

Dalton’s Atomic Theory (1808)

❑ Atoms cannot be subdivided, created, or destroyed❑ Atoms of different elements combine in simple whole-number ratios to form chemical compounds❑ In chemical reactions, atoms are combined, separated, or rearranged

❑ All matter is composed of extremely small particles called atoms❑ Atoms of a given element are identical in size, mass, and other properties; atoms of different elements differ in size, mass, and other propertiesJohn

Dalton

Modern Atomic TheorySeveral changes have been made to Dalton’s theory.Dalton said:

Atoms of a given element are identical in size, mass, and other properties; atoms of different elements differ in size, mass, and other properties

Modern theory states:

Atoms of an element have a characteristic average mass which is unique to that element.

Modern Atomic Theory #2

Dalton said:

Modern theory states:

Atoms cannot be subdivided, created, or destroyed

Atoms cannot be subdivided, created, or destroyed in ordinary chemical reactions. However, these changes CAN occur in nuclear reactions

Discovery of the ElectronIn 1897, J.J. Thomson used a cathode ray tube to deduce the presence of a negatively charged particle.

Cathode ray tubes pass electricity through a gas that is contained at a very low pressure.

Thomson’s Atomic Model

Thomson believed that the electrons were like plums embedded in a positively charged “pudding,” thus it was called the “plum pudding” model.

J.J. Thomson

Mass of the Electron

1909 – Robert Millikan determines the mass of the electron.

The oil drop apparatus

Mass of the electron is 9.109 x 10-31 kg

Conclusions from the Study of the Electron

❑ Cathode rays have identical properties regardless of the element used to produce them. All elements must contain identically charged electrons.❑Atoms are neutral, so there must be positive particles in the atom to balance the negative charge of the electrons❑ Electrons have so little mass that atoms must contain other particles that account for most of the mass

Rutherford’s Gold Foil Experiment

❑ Alpha particles are helium nuclei❑ Particles were fired at a thin sheet of gold foil❑ Particle hits on the detecting screen (film) are recorded

Try it Yourself!In the following pictures, there is a target hidden by a cloud. To figure out the shape of the target, we shot some beams into the cloud and recorded where the beams came out. Can you figure out the shape of the target?

The Answers

Target #1 Target #2

Rutherford’s Findings

❑ The nucleus is small❑ The nucleus is dense❑ The nucleus is positively charged

❑ Most of the particles passed right through❑ A few particles were deflected❑ VERY FEW were greatly deflected

“Like howitzer shells bouncing off of tissue paper!”

Conclusions:

Atomic Particles

The Atomic Scale

▪ Most of the mass of the atom is in the nucleus (protons and neutrons)▪ Electrons are found outside of the nucleus (the electron cloud)▪ Most of the volume of the atom is empty space

Helium-4

Image: User Yzmo Wikimedia Commons.

Atomic NumberAtomic number (Z) of an element is the number of protons in the nucleus of each atom of that element. This identifies the atom.

Mass NumberMass number is the number of protons and neutrons in the nucleus of an isotope.

Mass # = p+ + n0

8 8 18

18Arseni

c75

33

75Phosphor

us15

31

16

Electrons mass are so much smaller than a proton and neutron that they don’tContribute much to the overall mass of the atom ,therefore they are notCounted in the mass number of the atom.

Periodic TableMetals – good conductors of heat and electricityNonmetals- nonconductorsMetalloids -semiconductorsPeriods-horizontal rowsFamilies/groups-vertical columns

SilverSulfur

Periodic tableGroup names- elements in the same group will react

similarly

1- alkali metals (except Hydrogen) -most reactive metals

2- alkaline earth metals

3-12- transitions metals

17- halogens- most reactive nonmetals

18- noble gases (inert or unreactive)

Main group elements- groups 1,2,13-18

Percent error

• Percentage error = accepted value – experimental value

accepted value

Unit 1.1

Chapters 2.3, 3.3

Crash Course: chapter 2

IsotopesIsotopes are atoms of the same element having different masses due to varying numbers of neutrons.

Atomic MassesAtomic mass is the average of all the naturally isotopes of that element.

Carbon = 12.011

Purpose of Mass Spectrometry▪ Produces spectra of masses from the molecules in a

sample of material, and fragments of the molecules.▪ Used to determine

▪ the elemental composition of a sample▪ the masses of particles and of molecules▪ potential chemical structures of molecules by

analyzing the fragments▪ the identity of unknown compounds by determining

mass and matching to known spectra▪ the isotopic composition of elements in a molecule

StagesThe ionizer converts some of the sample into ions.

Mass analyzers separate the ions according to their mass-to-charge ratio. 

The detector records either the charge induced or the current produced when an ion passes by or hits a surface

Interpreting Mass Spectra

The height of each peak is proportional to the amount of each isotope present (i.e. it’s relative abundance). The m/z ratio for each peak is found from the accelerating voltage for each peak. Many ions have a +1 charge so that the m/z ratio is numerically equal to mass of the ion.

Calculating the relative atomic mass from mass spectrometry

1. Measure the height of each peak.

2. Calculate the percentage relative abundance

% abundance = amount of isotope x 100

total amount of all isotopes

3. Calculate the average mass mass of isotope A x abundance + mass of isotope B x

abundance

100

Atomic and molecular weights

Average atomic mass is the weighted average of all isotopes of an atom

H = 1.0078 amu amu=atomic mass unit

O = 15.9949 amu

Average atomic mass is on the periodic table and takes into account all isotopes of an atom and their abundance

Unit 1.2

Chapter 3.4

Crash course: chapter 2

The Mole (mol)1 dozen =

1 gross =

1 ream =

1 mole =

12

144

5006.02 x 1023

There are exactly 12 grams of carbon-12 in one mole of carbon-12.

Avogadro’s Number6.02 x 1023 is called “Avogadro’s Number” in honor of the Italian chemist Amadeo Avogadro (1776-1855).

Amadeo Avogadro

I didn’t discover it. Its just named after me!

Molar mass1 mol of book and 1 mol of feather

– Same number of items, NOT the same mass

Molar mass: mass in 1 mole of a substance (g/mol or g mol-1)– Molar mass = formula mass or average atomic

mass

1 mol a carbon = 6.02 x 10 23 atoms = 12.01 g

1 mol NaOH = 6.02 x 10 23 formulas = 40 g

1 mol CO = 6.02 x 10 23 molecules = 28 g

Calculations with Moles:Converting grams to moles

How many moles of lithium are in 18.2 grams of lithium?

18.2 g Li= mol Li

6.94 g Li

1 mol Li 2.62

Calculations with Moles:Converting moles to grams

How many grams of lithium are in 3.50 moles of lithium?

3.50 mol Li= g Li

1 mol Li6.94 g Li 24.3

Calculations with Moles:Using Avogadro’s Number

How many atoms of lithium are in 3.50 moles of lithium?

3.50 mol atoms1 mol

6.02 x 1023 atoms= 2.11 x 1024

Calculations with Moles:Using Avogadro’s Number

How many atoms of lithium are in 18.2 g of lithium?

18.2 g Li

= atoms Li

1 mol Li 6.022 x 1023 atoms Li

1.58 x 1024

6.94 g Li 1 mol Li

(18.2)(6.022 x 1023)/6.94

Moles

• Moles will convert to grams using molar mass

• Moles will convert to atoms/molecules/compounds using Avogadro's number

Calculating Formula MassCalculate the formula mass of carbon dioxide, CO2.

12.01 g + 2(16.00 g) = 44.01 g

Formula mass = mass in one chemical formula

Add the mass of all atoms in the formula

Calculating Formula MassCalculate the formula mass of carbon dioxide, CO2.

12.01 g + 2(16.00 g) = 44.01 g

Formula mass = mass in one chemical formula

Add the mass of all atoms in the formula

Conversions

• How many grams are in 2.50 mol of oxygen gas?

• (2.50 mol) (32 g O2/ 1mole)

• 80.0 g

• How many molecules are in 25.0 g of sulfuric acid?

• (25.0 g) (1 mol/98.1 g H2SO

4) (6.02x1023 molecules/1

mol)

• 1.53 x 10 23 molecules

• How many hydrogen atoms are in 25.0 g of sulfuric acid

• 2(1.53 x 10 23 molecules)

Ibuprofen, C13

H18

O2 has a molar mass of 206.29 g/mol.

If a bottle of ibuprofen contains 33 g of it, how many moles of ibuprofen are in the bottle and how many molecules are there?

0.16 moles, 9.6 x 1022 molecules

Unit 1.3

Chapter 3.5

Crash course: chapter 2

Percent composition

• We can find the mass percent of each element in a compound.

mass of element in compound x 100

total mass of compound

Calculating Percentage Composition

Calculate the percentage composition of magnesium carbonate, MgCO3.

Formula mass of magnesium carbonate:24.31 g + 12.01 g + 3(16.00 g) = 84.32 g

100.00

Formulas

❑ molecular formula = (empirical formula)n [n =

integer]

❑ molecular formula = C6H

6 = (CH)

6

❑ empirical formula = CH

Empirical formula: the lowest whole number ratio of atoms in a compound.

Molecular formula: the true number of atoms of each element in the formula of a compound.

Formulas (continued)

Formulas for ionic compounds (metals bonded to nonmetals) are ALWAYS empirical (lowest whole number ratio).Examples:

NaCl

MgCl2 Al2(SO4)3 K2CO3

Ionic compounds to not form molecules (they form crystals) so the formula doesn’t show the exact number of atoms in the compound but instead a ratio of how they bond.

Formulas (continued)

Formulas for molecular compounds (nonmetals bonded to nonmetals to form molecules) MIGHT be empirical (lowest whole number ratio).

Molecular:

H2O

C6H12O6 C12H22O11

Empirical:

H2O

CH2O C12H22O11

Empirical Formula Determination

1. Base calculation on 100 grams of compound.

2. Determine moles of each element in 100 grams of compound.

3. Divide each value of moles by the smallest of the values.

4. Multiply each number by an integer to obtain all whole numbers.

73.9% Hg and 26.1% Cl

73.9 g Hg and 26.1 g Cl

73.9/200.6 = 0.368 mol Hg

26.1/35.5 = 0.735 mol Cl

0.735/0.368 = 1.99

0.368/0.368 = 1

HgCl2

Make sure to go at least two places past decimal when finding molar mass

If you get a 1.5:1 ratio, double both 3:2

If you get a 1.33 : 1 ratio, triple both 4:3

What about 1.25: 1 ?

Empirical Formula Determination

Adipic acid contains 49.32% C, 43.84% O, and 6.85% H by mass. What is the empirical formula of adipic acid?

Empirical Formula Determination(part 2)

Divide each value of moles by the smallest of the values.

Carbon:

Hydrogen:

Oxygen:

Empirical Formula Determination(part 3)

Multiply each number by an integer to obtain all whole numbers.

Carbon: 1.50 Hydrogen: 2.50 Oxygen: 1.00x 2 x 2 x 2

3 5 2

Empirical formula: C3H5O2

Finding the Molecular FormulaThe empirical formula for adipic acid is C3H5O2. The molecular mass of adipic acid is 146 g/mol. What is the molecular formula of adipic acid?

1. Find the formula mass of C3H5O2

3(12.01 g) + 5(1.01) + 2(16.00) = 73.08 g

Finding the Molecular FormulaThe empirical formula for adipic acid is C3H5O2. The molecular mass of adipic acid is 146 g/mol. What is the molecular formula of adipic acid?

3(12.01 g) + 5(1.01) + 2(16.00) = 73.08 g

2. Divide the molecular mass by the mass given by the emipirical formula.

Finding the Molecular FormulaThe empirical formula for adipic acid is C3H5O2. The molecular mass of adipic acid is 146 g/mol. What is the molecular formula of adipic acid?

3(12.01 g) + 5(1.01) + 2(16.00) = 73.08 g

3. Multiply the empirical formula by this number to get the molecular formula.

(C3H5O2) x 2 = C6H10O4

example

• Empirical formula C3H

4

Molecular mass = 121 amu

What’s the molecular formula?

Mass of empirical formula = 40 amu

121/40 = 3.02

Molecular formula: C9H

12