Matter and Change Chapter 1.

Matter and Change Chapter 1

anything that has mass and takes up space
Matter anything that has mass and takes up space pure substances – compounds and elements mixtures – two or more pure substances mixed in the same container…not bonded

Solid, Liquid, or Gas?

SHAPE definite shape indefinite shape VOLUME definite volume indefinite volume Solid Liquid Gas

Physical Properties physical property – a characteristic that can be observed or measured without changing the identity of the substance state of matter melting point mass color

Physical Changes physical change – a change in a substance that does not involve a change in the identity of the substance breaking cutting dissolving boiling tearing

All phase changes are physical changes.

Chemical Properties chemical property – a characteristic that can be observed or measured with a change in the identity of the substance flammability reacts with an acid reacts with oxygen

Chemical Changes chemical change – a change in a substance that does involve a change in the identity of the substance color change gas released (often with an odor) energy change (light, heat, …) precipitate formed …four good indications of a chemical change.

Precipitate aqueous – dissolved in water
precipitate – a solid formed from two aqueous solutions during a chemical reaction

Heterogeneous Mixtures
Homogeneous Mixtures the same throughout…each sample contains the same ratio of ingredients Heterogeneous Mixtures different throughout…each sample contains a different ratio of ingredients

Separation of Mixtures
density PHYSICAL changes only magnetism by hand evaporation filtration chromatography

Separation of elemental Fe
magnetism This doesn’t work with iron that is bound into a compound, only with elemental Fe.

Distillation Apparatus

Homogeneous Mixture, Heterogeneous Mixture or Pure Substance?

Observations vs. Interpretations
observations – the facts…what you see or measure Example: The solution turned cloudy white. The test tube felt warm to the touch. interpretations – your opinion of what you see Example: From the previous observations, a chemical change has occurred.

Qualitative vs. Quantitative Observations
Qualitative observations – ones that do not involve numbers Quantitative observations – ones that DO use numbers (quantity)

Qualitative vs. Quantitative Which one is it?
1. A white precipitate was formed. 1. qualitative 2. The reaction produced 40.0 g of water. 2. quantitative 3. The test tube felt warm. 3. qualitative 4. The temperature rose 25 degrees. 4. quantitative

Extensive vs. Intensive Properties
extensive – DOES depend on the amount of matter present Example: mass, volume, amount of energy within a substance intensive – does NOT depend on the amount of matter present Examples: color, melting point, density, luster If mass and volume are extensive properties, why is density an intensive property?

Measurements and Calculations Pages 28-65
Chapter 2 Measurements and Calculations Pages 28-65

SI Prefixes Prefix Symbol Exponential tera T 1012 giga G 109 mega M
106 kilo k 103 hecto h 102 deka D or da 101 ----- 100 deci d 10-1 centi c 10-2 milli m 10-3 micro 10-6 nano n 10-9 pico p 10-12

SI Units Quantity Symbol Unit Name Unit abbreviation Length l meter m
Mass kilogram kg Time t second s Temperature T kelvin K Amount of a Substance n mole mol Electric Current I ampere A Luminous Intensity Iv candela cd

Size of Units 1L=1dm3 1cm3 = 1mL

Density Mass Density Volume Density is an INTENSIVE property… it does NOT depend on the amount of matter you have.

1. What is the density of a block of marble that has the dimensions 5
1. What is the density of a block of marble that has the dimensions 5.00cm length, 4.00cm width, and 15.5cm height and has a mass of 853 g?

2. Diamond has a density of 3. 26 g/cm3
2. Diamond has a density of 3.26 g/cm3. What is the mass of a diamond that has a volume of cm3?

3. What is the volume of a sample of liquid mercury that has a mass of 76.2 g, given that the density of mercury is 13.6 g/mL?

Significant Figures digit or figure – 0,1,2,3,4,5,6,7,8,9
significant digit or figure – a digit that helps you to understand the details of the entire number given to you

Significant Figures 5005 m L 4.500 g 100 cm3 100. cm3 g

Scientific Notation 50050 m = 5.005 x 104 m 0.00450 L = 4.50 x 10-3 L
Move the decimal until one non-zero digit appears to the left of the decimal. Be sure the power of the ten reflects the direction of that move. Keep the same number of significant figures in the scientific notation as your original. 50050 m = x 104 m L = 4.50 x 10-3 L

Accuracy (getting all measurements right) Precision (getting all measurements the same)

Percent Error What is the percent error for a mass measurement of 17.7 g, given that the correct value is 21.2 g?

In order to prepare for the chemistry exam on Wednesday, you should study for at least 1.0 hour the night before the exam. How many milliseconds is this? 3.6 x 106 ms

The summer Olympics showed us that Usain Bolt could run 10
The summer Olympics showed us that Usain Bolt could run meters per second. How fast is this in miles per hour? (3.281 ft = meters) 23.18 miles/hr

The volume of the Powerade® that I drank yesterday was 946 mL
The volume of the Powerade® that I drank yesterday was 946 mL. What is that volume in dm3? .946 dm3 or 9.46 x 10-1 dm3

The Building Blocks of Matter
Chapter 3: Atoms The Building Blocks of Matter An atom is the smallest particle of an element that retains the chemical properties of that element.

The Early Atom As early as 400 B.C., Democritus called nature’s basic particle the “atomon” based on the Greek word meaning “indivisible”. Aristotle succeeded Democritus and did not believe in atoms. Instead, he thought that all matter was continuous. It was his theory that was accepted for the next 2000 years. (Read page 43 of your textbook.)

Basic Laws of Matter Law of Conservation of Mass- mass is neither created nor destroyed during ordinary chemical reactions or physical changes. CH4 + 2O2 → 2H2O + CO2 16g + 64g → 36g g Antoine Lavoisier stated this about 1785

Basic Laws of Matter Law of Definite Proportions – no matter how much salt you have, it is always 39.34% Na and 60.66% Cl by mass. Example: Sodium chloride always contains 39.34% Na and 60.66% Cl by mass. 2NaCl → 2Na Cl2 100g → 39.34g g 116.88g → ? ? Joseph Louis Proust stated this in 1794.

John Dalton proposed this in 1803.
Basic Laws of Matter Law of Multiple Proportions- Two or more elements can combine to form different compounds in whole-number ratios. Example John Dalton proposed this in 1803.

Dalton’s Atomic Theory
In 1808, Dalton proposed a theory to summarize and explain the laws of conservation of mass, definite proportions, & multiple proportions. I was a school teacher at the age of 12!

Dalton’s Atomic Theory
John Dalton All matter is composed of extremely small particles called atoms. 2. Atoms of a given element are identical in size, mass, and other properties.** 3. Atoms cannot be subdivided, created, or destroyed.** 4. Atoms of different elements combine in simple whole-number ratios to form chemical compounds. 5. In chemical reactions, atoms are combined, separated, or rearranged. **Today, we know these parts to have flaws.

Flaws of Dalton’s Theory…
2. Atoms of a given element are identical in size, mass, and other properties. 3. Atoms cannot be subdivided, created, or destroyed. Isotopes – atoms with the same number of protons but a different number of neutrons Subatomic particles – electrons, protons, neutrons, and more

The Atom Atom - the smallest particle of an element that retains the chemical properties of that element. CARBON

Subatomic Particles Protons- positively charged particles found in the nucleus of an atom. Neutrons- neutral particles found in the nucleus of an atom. Electrons- negatively charged particles found in the electron cloud.

Joseph John Thomson In 1897 the English physicist Joseph John Thomson was able to measure the ratio of charge of the cathode ray particles to their mass. He found that the ratio was always the same regardless of the metal used to make the cathode or the nature of the gas inside the cathode ray tube. Thomson concluded that cathode rays were composed of identical, negatively charged particles called electrons.

Cathode Ray Tube Experiment Accomplishments
Proved that the atom was divisible and that all atoms contain electrons. This contradicted Dalton’s Atomic Theory. This allowed a new model of the atom.

Discovery of the Nucleus
In 1911, Ernest Rutherford performed a Gold Foil Experiment. He and his colleagues bombarded a thin piece of gold foil with fast moving, positively charged alpha particles.

Alpha Particles Alpha (a) particles are Helium-4 nuclei.
This means they are two protons and two neutrons (with no electrons). Thus, they are positive.

Gold Foil Experiment

Gold Foil Experiment The volume of the nucleus was very small compared to the volume of the atom. Therefore, most of the atom was composed of empty space. Niels Bohr later found that this empty space was where the electrons were located.

Bohr’s Model of The Atom

Atomic Number atomic number (Z) - the number of protons in the nucleus of each atom of a given element. – (Henry Moseley) The number of p+ identifies the element. Atomic Number increases from left to right on the periodic table.

Electrons The number of electrons in a neutral atom is equal to the number of protons in that atom. e- = p+ Electrons can be lost or gained. When electrons are lost or gained, ions are formed.

Metals vs. Nonmetals Metals form cations. Na Na+ + 1e-
Nonmetals form anions. Cl + 1 e Cl-

# of neutrons = mass number – atomic number
mass number (A)- the number of p+ & no in the nucleus of an atom. # of neutrons = mass number – atomic number Why aren’t electrons included when determining the mass number of the atom?

Isotopes isotope- two or more atoms having the same atomic number (same #p+) , but different mass numbers (due to different #no).

Isotope Notation Nuclear Notation Hyphen Notation
Uses the elements symbol followed by a hyphen & the mass number. C-12

How many protons, neutrons & electrons are there in the following?
Cl Cl-1 Br S-2 N Fe+3

Changes in the Nucleus Nuclear Reaction- changes that occur in the atom’s nucleus. Nuclear reactions can change the composition of an atom’s nucleus permanently.

Types of Radioactive Decay
Alpha Radiation (a)- stream of high energy alpha particles. Consists of 2 protons & 2 neutrons making it identical to a He-4 nucleus. Alpha particles can be represented by: Most alpha particles are able to travel only a few centimeters through air and are easily stopped by clothing etc.

Alpha Decay parent daughter

Types of Radioactive Decay
Beta Radiation (b) – consists of a stream of high speed electrons. These electrons are not electrons that are in motion around the atom’s nucleus. Beta particles can be represented by: Can penetrate through clothing and damage skin.

Beta Decay parent daughter

Checking for Understanding
alpha decay beta decay

The Mole mole (mol)- SI Unit for the amount of a substance that contains as many particles as there are atoms in exactly 12g of carbon-12. A unit of counting, like the dozen.

1 mole = 6.0221415 X 1023 1 mol = 6.02 x 1023 Avogadro’s Number
Avogadro’s Number - the number of particles in exactly one mole of a pure substance. 1 mole = X 1023 1 mol = 6.02 x 1023 Amedeo Avogadro

Atomic Mass atomic mass - the mass of one mole of an atom
Atomic mass is expressed in atomic mass units (amu) or (u) or g/mol. Can be found on the periodic table. All atomic masses are based on the atomic mass of carbon-12 being 12 amu.

Molar Mass molar mass - the mass of one mole of a pure substance. Molar mass is written in units of amu or g/mol.

Atomic mass vs. Molar mass
atomic mass - the mass of one mole of an atom. molar mass - the mass of one mole of a pure substance.

Atomic Mass vs. Molar Mass
Example Atomic Mass Molar Mass Na Ag C O 22.99 g/mol 22.99 g/mol g/mol g/mol 12.01 g/mol 12.01 g/mol 16.00 g/mol 16.00 g/mol

Molar Mass of Compounds
H2O C6H12O6 NaCl Cl2 (NH4)3PO4 CuSO4·5H2O 18.02 g/mol g/mol 58.44 g/mol 70.90 g/mol g/mol g/mol

The Mole Bridge

Cheer I say grams, you say molar mass. grams – molar mass

Grams to Moles Converting grams to moles: divide by molar mass.
1. How many moles of Ca are in 5.00g of Ca?

Moles to Grams Converting from moles to grams: multiply by molar mass
What is the mass in grams of 2.25 moles of Fe?

Types of Particles 1 mole = 6.02 x 1023 particles Atoms – C, Cu, He
Molecules – O2, C12H22O11, CO2 (all nonmetals in the formula) Formula units – NaCl, CaCl2, Mg(NO3)2 (includes a metal in the formula) 1 mole = 6.02 x 1023 particles

Particles to Moles Converting particles to moles: divide by Avogadro's Number. How many moles of Pb are in 1.50 X 1025 atoms of Pb? 2.49 x 101 moles Pb

Moles to Particles Converting moles to atoms: multiply by Avogadro's Number. How many molecules of NO are in 0.87 moles of NO? 5.2 x 1023 molecules NO

Grams to Moles to Particles
Example: How many molecules of N2 are in 57.1g of N2?

Particles to Moles to Grams
Example: How many grams of NaF are in 7.89 X 1024 formula units of NaF?

Arrangement of Electrons in Atoms
Chapter 4 Arrangement of Electrons in Atoms The emission of light is fundamentally related to the behavior of electrons.

The Concept of Energy Energy – the ability (capacity) to do work
Units of Energy: calorie (cal), Joule (J) Work is done when an object moves some distance in response to a force (pushing or pulling). work = force x distance

Forms of Energy Mechanical energy Chemical energy Solar energy
Radiant or Electromagnetic energy Electrical energy Nuclear energy Thermal energy Sound energy Magnetic energy

Where Does Energy Go? Law of Conservation of Energy – energy cannot be created nor destroyed Energy is continually transferred from one thing to another, never disappearing.

PE – “energy of position”
Types of Energy kinetic energy – the energy that objects have because they are moving KE – “energy of motion” potential energy – the energy that is available for doing work at some later time PE – “energy of position”

Energy Units What units are used for energy? joule = J (SI unit)
calorie = cal (non SI unit) Calorie = Cal (food calorie = 1000 cal) 1.00 cal = J Example: How many calories of energy is 1067 joules? James Prescott Joule

Measurable Properties of Light
Wavelength (l) – the distance between corresponding points on adjacent waves Units: meters, nanometers, etc. 1m = 1x109nm

Measurable Properties of Light
Frequency (u)- number of waves that pass a given point in a specific time (usually one second) Unit: Hertz (Hz)

Types of Electromagnetic Radiation
Gamma rays X-rays Ultraviolet Visible Infrared Microwaves Radio waves High energy, E High frequency, u Violet Indigo Blue Green Yellow Orange Red High wavelength, l Memorize these in order. Know E, u, and l order.

Relationship between u and l
c = lu c = speed of light in m/s l = wavelength in m u = frequency in 1/s

What is the constant, h? E = energy in J h = constant in J.s
u = frequency in 1/s or Hz

Combining the Equations you get…
l u = E = h u l

Check for Understanding
13. What would the energy be for light with a frequency of 5.68 x /s? 14. What is the energy of light that has a wavelength of 7.89 x m?

Check for Understanding

Photoelectric Effect Photoelectric effect- the emission of electrons from a metal when light shines on the metal. The photoelectric effect does not occur when the light’s frequency is below a certain amount regardless of the intensity of the light.

Absorption vs. Emission
Absorption – when energy is “taken in” by electrons Emission – when energy is “given off” by electrons

Energy States of an Electron
Ground State- the lowest energy state of an atom. (stable) Excited state- state in which an atom’s potential energy is increased from that of the ground state. (unstable).

Spectroscopy Spectroscope – used to separate light into a spectrum by wavelength so it can be examined. Line-Emission Spectrum – produced when an electron jumps from a higher energy level to a lower energy level. Acts as an atomic fingerprint.

How to use a Spectroscope

Discovery of the Line-Emission Spectrum
Hydrogen atoms were excited by passing a high-voltage current through hydrogen gas causing the gas to glow a lavender color. When viewed with a spectroscope (diffraction grating or prism) the lavender light separated into four narrow lines of different color.

Hydrogen Line Spectrum
Each of the lines seen in the hydrogen spectrum is a result of light at a different wavelength. Since light of a particular wavelength has a definite frequency and a definite energy, the lines of the hydrogen spectrum must be a result of the emission of photons with specific energies.

Other Emission Spectrum
helium carbon

The Bohr Model

Bohr Model of the Atom The fact that hydrogen atoms only released photons of specific frequencies indicated that differences between the atom’s energy levels were fixed. The puzzle behind the hydrogen-atom spectra was solved in 1913 by Niels Bohr.

Bohr Model of the Atom The orbits are separated from one another by empty space where the electrons cannot exist. The energy of the electrons becomes higher as they get farther away from the nucleus. Electrons can move from one energy level to the next by gaining or losing a finite amount of energy.

Arrangement of Electrons in Atoms
Chapter 4 – Part II Arrangement of Electrons in Atoms The emission of light is fundamentally related to the behavior of electrons.

Heisenberg Uncertainty Principle
In 1927, Werner Heisenberg stated that it is impossible to simultaneously determine both the position and velocity of an electron or any other particle. This became known as the Heisenberg Uncertainty Principle. Quantum theory was more widely accepted after this proposal.

Quantum Model of the Atom
The dots represent 1 electron and the region in which you are likely to find it 90% of the time.

Energy Levels Quantum numbers describe the location of electrons in an atom.

Principal Quantum Number, n
principal quantum number, n – indicates the main energy level occupied by the electron n = 1→7 (notice how many periods are on the periodic table…7 !) (Does this sound like the Bohr model?)

Energy Sublevels Every energy level can be broken down into sublevels.
No more Bohr model Every energy level can be broken down into sublevels. Each sublevel has an energy that is slightly different. Not the traditional step ladder anymore, but still you can’t have energies in between the rungs of our new ladder.

Orbitals orbital – a 3-D region in space within an energy level in which the electron is most likely to be found (90% of the time). An orbital can hold a maximum of 2 e- ! Pauli Exclusion Principle

Sublevels in the Quantum Model
s orbital 2 electrons 1 orientation

Sub Levels in the Quantum Model
p orbitals: px, py, pz 6 electrons total 3 orientations

d orbitals: dxy, dxy, dyz, dx2-y2, dz2 10 electrons total 5 orientations

f orbitals 14 electrons total 7 orientations

Summary of orbitals energy levels sublevels n = 1 1s n = 2 2s, 2p
n = 3 3s, 3p, 3d n = 4 4s, 4p, 4d, 4f n = 5 5s, 5p, 5d, 5f n = 6 6s, 6p, 6d n = 7 7s, 7p

Pauli Exclusion Principle – a maximum of 2 e- may occupy one orbital, both with opposites spins
Orbital Diagrams Aufbau Principle – fill in the lowest possible energy orbital Aufbau Principle – fill in the lowest possible energy orbital Hund’s Rule – within equal energy orbitals, the e- are distributed to have the maxiumum unpaired e- possible 3d 4s Pauli Exclusion Principle – a maximum of 2 e- may occupy one orbital, both with opposite spins Aufbau Principle – fill in the lowest possible energy orbital 3p 3s 2p 2s 1s

Valence Eletrons Br: 1s22s22p63s23p64s23d104p5
valence electrons – electrons in the highest energy level Br: 1s22s22p63s23p64s23d104p5 1st E level 2nd E level 3rd E level 4th E level 4th is the highest Energy level, so there are 7 valence electrons

Chapter 5: Periodic Trends
Who told the elements where to go? MENDELEEV! (Father of the periodic table)

History of the Periodic Table Early 1800’s
J.W. Dobereiner organized elements into triads. Triads- groups of three elements with similar properties. The middle element of the triad was thought to be an approximate average of the properties of the first & third element.

History of the Periodic Table Mid 1800’s
J.A.R. Newlands- developed the law of octaves in 1865. Law of octaves- stated that if elements were arranged by increasing atomic mass, the properties of the eighth element were similar to properties of the first, the ninth like those of the second, the tenth like those of the third.

History of the Periodic Table Mid 1800’s
In 1869, Dmitri Mendeleev published the first periodic table. Mendeleev arranged the elements horizontally by increasing atomic mass and placed elements in groups (vertically) based on similar properties.

History of the Periodic Table Early 1900’s
In 1913, Henry Moseley developed the Modern Periodic Table. He determined the atomic numbers of many elements and then used those values to arrange the periodic table in rows by increasing atomic number and in columns by similar properties.

Metals, Nonmetals, & Metalloids
Make up most of the periodic table. Ductile- can be drawn into wire Malleable- can be hammered into thin sheets. Lustrous- shiny Good Conductors of heat & electricity Located to the left of the step ladder on the periodic table.

Brittle-break when hammered. Poor conductors of heat & electricity. Lack luster Located to the right of the step ladder on the periodic table.

Semimetals. Properties of both metals & nonmetals. Located along the step ladder on the periodic table. Examples: B, Si, Ge, As, Sb, Te

Alkali Metals- Group 1 Most reactive group of metals.
Usually found in combined form as a salt due to their high reactivity. Combine vigorously with nonmetals especially groups 16 & 17. React readily with water. Soft and silvery appearance.

Alkaline Earth Metals- Group 2
Found in the earth’s crust but not in the elemental form due to their high reactivity. They are usually found in rock structures. 2nd most reactive group of metals. More dense than group 1. Shiny silvery-white color.

Transition Metals- Groups 3-12
Compose the d-block. Typical metallic properties. Good conductors. Lustrous. Produce colored ions.

Main Group Elements Groups 13-18
Compose the p-block. Properties of elements vary greatly. Contains all of the nonmetals & metalloids as well as some metals.

Halogens -Group 17 Most reactive groups of non-metals.
React vigorously with metals (especially groups 1 & 2) to produce salts. Fluorine is a poisonous pale yellow gas, chlorine is a poisonous pale green gas, bromine is a toxic and caustic brown volatile liquid, and iodine is a shiny black solid which easily sublimes to form a violet vapor on heating. Found in nature in the combined state.

Noble Gases- Group 18 Least reactive of all elements.
Often called inert gases. All are gases. The noble gases are all found in minute quantities in the atmosphere, and are isolated by fractional distillation of liquid air. Helium can be obtained from natural gas wells where it has accumulated as a result of radioactive decay.

Inner Transition Metals- Periods 6 & 7
Compose the f-block. Fill in Between Groups 3 & 4 on the Periodic Table. Lanthanides (Period 6)- Rare Earth Metals Shiny reactive metals Actinides (Period 7) Unstable & radioactive metals. Most are laboratory made.

Reactivity of Metals Trend
Period Trend- Metals increase in activity from right to left on the periodic table. The alkali metals are the most reactive group of metals. Group Trend- Metals increase in reactivity from top to bottom with a group. Ra is the most reactive alkaline earth metal.

Reactivity of Nonmetals Trend
Period Trend- Nonmetals increase in activity from left to right on the periodic table with the exception of the noble gases. The halogens are the most reactive group of nonmetals. Group Trend- Nonmetals increase in reactivity from bottom to top with a group. F is the most reactive halogen.

Reactivity Trends What is the most reactive metal on the periodic table? Explain. Circle the most reactive nonmetal in each row: Te Po S Br I Cl

Atomic Radii Atomic Radius- ½ the distance between the nuclei of two identical atoms that are bonded together. Measure from nucleus to nucleus and divide by two. The larger the atomic radius the larger the atom.

Atomic Radii Trend Atomic Radii increases as you travel down & left on the periodic table!

Atomic Radii Trend

Atomic Radii Trends 1. Circle the atom with the largest atomic radii.
a. I F Cl b. O B N c. S Sb Sr 2. Put the following elements in order of increasing atomic radii. a. Na, P, Rb 3. Which element has the largest atomic radii?

Ionic Radii Cation- a positive ion formed from the loss of electrons. Ex. Ca+2 Metals typically form cations. Cations are smaller than their parent atoms. Example: Na > Na+

Ionic Radii Anion- a negative ion formed from the gain of electrons. Ex. Cl-1 Nonmetals typically form anions. Anions are larger than their parent atoms. Example: Cl-1 > Cl

Ionic Radii Trends 1. Circle the larger ion. a. Cl-1 Cl b. Na Na+1
c. N N-3 d. Fe+2 Fe+3 e. C+4 C-4

Ionization Energy Ionization Energy- the energy required to remove an electron from an atom. Nonmetals have higher ionization energies than metals. The lower the ionization energy of an atom the easier it is to remove an electron.

Ionization Energy Trend
General Group Trend (representative elements)- Ionization energy increases as you travel up a group in the periodic table. Reason: As you travel up a group there are less principle energy levels, therefore, electrons are located closer to the nucleus. It is harder to remove electrons that are closer to the nucleus to due the nuclear charge attraction.

Ionization Energy Trend
Ionization energy increases as you travel up & right on the periodic table!

Ionization Energy Trends
1. Circle the atom with the highest ionization energy. a. F O B b. As N P c. F C Li

Electron Affinity Electron Affinity- energy change that occurs when an electron is acquired by a neutral atom. Nonmetals have a more negative (LOWER) electron affinity than metals. Metals typically have a more positive (HIGHER) electron affinity than nonmetals. Noble gases have an electron affinity between metals & nonmetals.

Electron Affinity Trends
1. Circle the atom with the highest electron affinity. a. N S Ar b. O Ne Ca c. Cu Ar Kr

Electronegativity Electronegativity- measure of the ability of an atom in a chemical compound to attract electrons. The higher the electronegativity the greater the attraction for electrons. The electronegativity scale was determined by Linus Pauling. The electronegativity scale ranges from 0 to 4. Since some noble gases do not form compounds they have electronegativity values of 0. Used to determine the polarity of an bond.

Electronegativity Trend
Electronegativity increases as you travel up & right on the periodic table!

Electronegativity Trends
1. Circle the atom with the highest electronegativity. a. F S Sn b. Li N O

Chemical Bonding Chapter 6 Pages

15-1 The Attachment Between Atoms
atoms combine to form ionic bonds covalent bonds (M + NM) (NM + NM) chemical bond – a mutual electrical attraction between the nuclei and valence electrons of two atoms that binds the atoms together

Ionic Bonding cation and anion
ionic bond – when electrons are taken by one atom from another atom metal and a nonmetal NaCl cation and anion (The charges are “hidden” to make a neutral compound.)

The simplest ratio of the packed ions is called:
The Formula Unit Ex: NaCl

Ions cations (+) anions (-)
monatomic ions – ions formed from one atom Examples: Na+ or O-2 polyatomic ions - ions formed from two or more atoms bonded together Examples: NH4+ or SO4-2

Ionic Compounds solid at room temperature
high melting points (thus are usually solid at RT) formula unit represents the lowest ratio of ions that combine to form a neutral compound most are crystalline solids when dissolved in water, the ionic compounds will break up into ions (dissociate) the solutions of ionic compounds will conduct electricity (electrolytes)

Covalent Bonding covalent bond – when electrons are shared between two atoms two nonmetals No ions formed! (no electrons are taken)

There is another type of bond, not purely covalent and not purely ionic.
Nonpolar Covalent Pure Covalent Polar Covalent Ionic

Electronegativity electronegativity – a measure of the ability of an atom in a chemical compound to attract electrons from another atom in the compound The difference in electronegativity values for two atoms will indicate whether the two atoms form an ionic bond or a polar or nonpolar covalent bond.

Electronegativity Differences
0.0 to 0.3 nonpolar covalent 0.3 to 1.7 polar covalent 1.7 and up ionic Learn these values…Although the general rule is a metal and nonmetal will form an ionic bond and two nonmetals will form a covalent bond.

Ionic, Polar Covalent, or Nonpolar Covalent?
What kind of bond would each pair form? N and S S and C Mg and Cl C and F Ba and O

Valence Shell Electron Pair Repulsion Theory VSEPR Theory
Repulsion between the sets of valence-level electrons surrounding an atom causes these sets to be oriented as far apart as possible.

Regions of Electron Density
What is a Region of electron density? Single bond (2e- connecting 2 atoms) Double bond(4e- connecting 2 atoms) Triple bond(6e- connecting 2 atoms) Lone pair (unbonded pair) (2e- alone on an atom)

LINEAR 180o sp hybridization 2 Regions of Electron Density 2 Bonds
bonded pair of electrons bonded pair of electrons

TRIGONAL PLANAR 120o sp2 hybridization 3 Regions of Electron Density 3 Bonds
3 bonded pairs of electrons

TETRAHEDRAL 109.5o sp3 hybridization 4 Regions of Electron Density 4 Bonds
4 bonded pairs of electrons

TRIGONAL PYRAMIDAL 107o sp3 hybridization 4 Regions of Electron Density 3 Bonds & 1 Lone Pair
1 lone pair of electrons 3 bonded pairs of electrons

BENT 105o sp3 hybridization 4 Regions of Electron Density 2 Bonds & 2 Lone Pairs

All of these have 4 regions of electron density and sp3 hybridization
(although the number of bonded pairs is different)

TRIGONAL BIPYRAMIDAL 120o & 90o 5 Regions of Electron Density 5 Bonds

OCTAHEDRAL 90o 6 Regions of Electron Density 6 Bonds
SF6

SQUARE PLANAR 90o 6 Regions of Electron Density 4 Bonds & 2 Lone Pairs
ICl4-

Hybridization Regions of e- Density Hybridization 2 sp 3 sp2 4 sp3

Chapter 6: Chemical Bonding
Part II – Polarity

What are the three types of bonds?
What do the differences in electronegativity indicate between two atoms? the bond type that they will form What are the three types of bonds? Ionic (difference of 1.7+) Polar covalent (difference of ) Nonpolar covalent (difference of )

Nonpolar Covalent Bond
Nonpolar covalent bond – a covalent bond in which the electrons are shared equally Example: Cl2 Both chlorines have the same electronegativity.

Polar Covalent Bond H Cl
Polar covalent bond – a covalent bond in which the electrons are shared unequally Example: HCl H Cl Which end represents the hydrogen end & which end represents the chlorine end? How do you know this answer? Cl has a higher electronegativity

Dipole dipole – partial negative or partial positive charge formed during unequal sharing of electrons (in polar bonds only) d+ d- The direction of a dipole is from the dipole’s positive pole to its negative pole.

Nonpolar vs. Polar Covalent Bond
There is no dipole drawn over the H2 molecule… WHY? The dipole is drawn over the HCl molecule. The arrow points towards the more electronegative element, Cl. There is not an uneven distribution of electrons.

Nonpolar vs. Polar Covalent Bond
The electrons are not shared equally within the HCl molecule. Compare the electron clouds for the H2 and HCl molecules. Why are they different? Do we usually draw the electron clouds with the Lewis structures? NO

Each dipole represents a polar covalent bond.
Example: PCl3 has three dipoles because there are three polar bonds. These three dipoles are drawn beside the bonds, pointing towards the more electronegative element, Cl. What is the more electronegative element, P or Cl? P Cl Cl Cl So, how would the dipoles be drawn? Each dipole represents a polar covalent bond.

You Try It! Draw the Lewis Structures on your white board for the following molecules (Aha!…they must all be covalent, then!) and LABEL ALL THE DIPOLES along each bond. CO2 HBr NH3 Draw them on your white board and only when you are ready to check the Lewis structures AND dipoles, go on to the next slide.

Answers Did you make CO2 linear? Did you make NH3 trigonal pyramidal?
Did you show the dipoles pointing to the more electronegative element?

Dipoles within Molecules
There are dipoles created along bonds, however, a molecule can exhibit overall polarity. Here is the dipole along the first H-O bond… An overall dipole is created. Here is the dipole along the second H-O bond…

An overall dipole is created.
Two individual dipoles along each bond cause the molecule to have a greater electron cloud towards the oxygen end leaving the other end of the molecule more positive. This causes a partial negative end towards the oxygen and the a partial positive end towards the hydrogens (you can only have two “ends” to a small molecule. An overall dipole is created.

The bond is polar and so is the overall molecule.
When you have a two atom molecule, it is easy to tell if it is a polar molecule. The bond is polar and so is the overall molecule. partially negative end partially positive end

Polarity of a Molecule Polarity of a molecule depends upon two things
the polarity of the bonds the shape of the molecule All molecules with only nonpolar bonds are nonpolar molecules. Molecules with polar bonds, may or may not be polar molecules, depending on the shape.

carbon monoxide, CO Is the bond polar? YES Is the molecule polar?
YES…a two atom molecule is polar if the bond is polar.

ammonia, NH3 Are the bonds polar? YES YES…the molecule is not symmetrical and does have a partial positive end and a partial negative end. You can draw on overall dipole. Is the molecule polar?

methane, CH4 Are the bonds polar? YES NO…the molecule is symmetrical and does NOT have a partial positive end and a partial negative end. (The center and the outside does not count as an “end”.) Is the molecule polar?

Polar or Nonpolar Molecule?
Rule #1: All molecules with nonpolar bonds are always nonpolar molecules. Rule #2: Molecules with polar bonds, will be polar only if they are nonsymmetrical (like NH3).

Draw them on your white board. Click to see the answers.
You Try It! #2 Draw the following molecules, including the dipoles along each bond. Determine whether they are polar or nonpolar molecules. CH3Cl SF6 O2 H2S KCl Draw them on your white board. Click to see the answers.

Formula Writing and Naming
Chapter 7 – Part I Formula Writing and Naming May I please have a glass of dihydrogen monoxide

Be sure to know how to name…
Ions polyatomic and monotomic, cations and anions Ionic compounds ones with and without Roman numerals Molecular/Covalent compounds use prefixes for these (mon-, di-, tri-, tetra-, …) Acids

Chapter 7 – Part II Math with Formulas

Percent Composition What is the percent composition of oxygen in carbon dioxide? Percent composition is determined by mass. You need to know the total mass of the element for the bottom of your fraction and the mass of the part of the compound in question for the top of the fraction. multiplied by 2 because there are 2 oxygen in the formula molar mass of oxygen molar mass of carbon dioxide

Empirical Formulas What is the empirical formula for glucose, C6H12O6?
An empirical formula is just the lowest ratio of the molecular formula. It is not representative of the real molecule, but it just shows the ratio of atoms in that molecule. C6H12O6 has a ratio of 6:12:6 which can be reduced, so the answer is empirical formula is CH2O (1:2:1 ratio)

Determining Molecular Formulas from Empirical Formulas
What is the molecular formula for a compound that weighs 110.g/mol and has an empirical formula of C3H3O? The molecular formula must weigh some multiple of the empirical formula. First determine the mass of the empirical formula. C3H3O weighs 55 g/mol and the real compound weighs 110. g/mol, or twice that much. So, the real formula is 2x C3H3O or C6H6O2

Matter and Change Chapter 1.

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Matter and Change Chapter 1.

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