High school science course 1 credit/unit with lab

High school science course 1 credit/unit with lab
Chemistry 1 High school science course 1 credit/unit with lab

Module 1: Measurements What is chemistry? PLANTS METALS HINT: CERAMICS
The study of matter…what is matter?…anything that takes up space and has mass…what is mass?...the number of atoms it has! WHAT WOULD NOT BE MATTER ON EARTH? PLANTS METALS HINT: CERAMICS POLYMERS ANIMALS

Units of Measurements It is essential to include all units used in an experiment!!! NASA example in book, p. 5 – this was a multi-billion dollar mistake!!! TEAM A ~ used Metric units TEAM B ~ used English units If you work a problem or experiment and the units are left off, your answer will be considered incorrect!!! MASS = how much space an object occupies + how much matter it has WEIGHT = how hard gravity is pulling on an object

The Metric System Prefixes:
MASS = how much space an object occupies + how much matter it has On the MOON or EARTH, mass would be the same WEIGHT = how hard gravity is pulling on an object On the MOON, weight is less due low gravity Basic unit of mass in the Metric System is the gram (g) = 1 g = 2,300 g = _____ kg 2.3 Prefixes: Used when GREATER or LESSER than a gram

Converting Between Units
= 11.1 cm = _____ meters Use the factor-label method: 1 cm = o.o1 m 11.1 cm 0.01 m 0.111 m 1 1 cm Given unit Conversion Factor Wanted unit = in = _______ cm 26.3 in 2.54 cm cm 1 1 in Given unit Conversion Factor Wanted unit

Derived Units These are units that come from the basic units of the metric system. Many units in Chemistry are derived. ie: Volume is cubic centimeters = cm3 Remember to measure container before measuring Object/Substance graduations meniscus Graduated Beaker Graduated Cylinder Mass Scale

Multiply and Dividing Derived Units
Units do NOT have to match, but must NOT be lost in the results 1.1 m m m2 unit of area 2 dimensions 1.1 cm cm cm cm³ unit of volume 3 dimensions cm³ is the usual unit for volume, and equals ml What is in cubic meters? 3.96 cm³ 3.96 cm³ 0.01 m m3 1 1 cm Given unit Conversion Factor Wanted unit

Accuracy, Precision, & Significant Figures
Accuracy : how close to the true value you can get, ie: in lab you obtain a mass of 3.2 kg. Actual weight was 10 kg. You are not very accurate. Precision: closeness of two or more measurements on a device , ie: in lab you obtain a mass of 3.2 kg x 3 times. It is NOT dependent on accuracy. The Analogy: You throw the dart toward the bull’s eye. You are close each time, then you accurate. If you miss the bull’s eye each time, but constantly hit the board near the same place, then you are not accurate, but you are precise!

We need to know how to evaluate the accuracy and precision of our measurements!
That’s where significant figures come in: They come from a measuring device. They are measured digits.

Scientific Notation Scientific Notation!!!
Sometimes it is difficult to write tiny or huge numbers with precision, so there is… Scientific Notation!!! We write numbers, then, with a decimal point, and then multiple by 10 with exponents. To write a large number in scientific notation, you move the decimal to the left and the exponent will have to be positive. 22 20,000,000,000,000,000,000,000 particles/breath of air = 2.0 x particles To write a small number (between 0 and 1) in scientific notation, you move the decimal to the right and the exponent will have to be negative. kg = mass of proton (in center of atom) = 1.67 x kg -27

Measuring Temperature
Fahrenheit; symbol: F = of or denoting a scale of temperature on which water freezes at 32° and boils at 212° under standard conditions. Celsius; symbol: C = of or denoting a scale of temperature in which water freezes at 0° and boils at ° under standard conditions. °C = 25 T(°F) = 9/5× T(°C) + 32 9 5 25 1 T(°F) = 9/5 x = x + 32 °F + 32 = = 77 °F = 50 T(°C) = 5/9 ( ) T(°F) 5 9 18 1 T(°C) = 5/9 (50 -32) = 5/9 (18) = x = 10 °C

°C °C = 100.0 °C Kelvin (K) = + 273.15 = 100 + 273.15 = 373.15 K

Module 2: Atoms & Molecules
-protons -neutrons -electrons Many atoms = Molecule

You cannot make or destroy matter. You can only change its form.
Law of Mass Conversion You cannot make or destroy matter. You can only change its form. Experiment: Burn a log of wood (Mass of log before burning = 100 g) Ashes Oxygen (mass =? g ) (mass = 73 g) 100 g = 73 g + Oxygen Oxygen = 100 g – 73 g = 27 g

Elements: Basic Building Blocks
Any substance that cannot be broken down into smaller components Na Fe Simple substances: (See Periodic Table) Cl He 92 come from nature + 26 manufactured in lab by scientists = 118 Total Each element on the Periodic Table has an abbreviation Some come from the LATIN language Others from the scientist who discovered them Others from the place it was discovered Notice the jagged line: separates metals and nonmetals - Exception: Hydrogen – it is a nonmetal

Compounds The Law of Definition Proportions (Proust’s Law)
Substances that consist of identical molecules The Law of Definition Proportions (Proust’s Law) Several elements come together in the same proportions Example: To make table salt Can also use the Law of Mass Conversions Example: 6.5g Na g Cl = 16.5 g NaCl 6.5g Na g Cl = 16.5 g NaCl, but what are the proportions for 25.0g of NaCl? First rewrite the equation to: 16.5g x U = 25g U = 25.0g/16.5g= 1.52 Next we need to increase our proportions by 1.5 6.5g x 1.52 = 9.88g Na maximum 10.0g x 1.52 = 15.22g Cl maximum So9.88g Na g Cl = 25g NaCl S

The Law of Multiple Proportions (Dalton’s Law)
When 2 elements combine to form more than one compound Includes 4 basic assumptions: All elements are made of tiny particles called atoms * All atoms of the same element have the same properties Atoms of different elements have different eproperties Compounds are formed when atoms join together in simple, whole-number ratios Example: CO2 has 1 carbon atom + 2 oxygen atoms There can not be 1.5 atoms *Atoms cannot be seen, so from now on we simply accept their existence without proof. Today we have found that 2 of Dalton’s assumptions are not correct: Atoms can be split apart There is the existence of isotopes

Molecules: Building Blocks of Compounds
All elements are composed of identical atoms Elements join together in groups called molecules Identical molecules create compounds Element Molecule Compound Mixture

Each element has a name and abbreviation
Each compound has a name and abbreviation, after the elements that make them up Now, we need to count the exact # of atoms for each element Finally, we can write its formula: NNaO3 He C Ni Br He Nitrogen Sodium Oxygen 1 1 3

Classifying Matter NaNO3
We have the formula…now we can name the compound! First, we will classify it as IONIC or COVALENT NaNO3 IONIC Dissolves in water Conducts electricity Made up of at least one metal atom and at least one nonmetal atom. NaHCO3 (Baking Soda) COVALENT Does not dissolve in water Cannot conduct electricity Made up of solely nonmetal atoms. 4. C12H22O11 (Table Sugar)

s s s N Metals Metalloids Nonmetals

Naming Compounds Ionic
We use a different system for Ionic or Covalent Compounds Ionic Start with the name of the first atom in the molecule. Replace the name of the next atom in the molecule with its –ide name. Putting those 2 names together gives the compound’s name. EXAMPLE: NaCl ~ 1 atom of sodium + 1 atom of Chlorine becomes chloride NAME: Sodium Chloride

Covalent Multiple combinations can occur
We use prefixes in front of each element We drop the prefixes on the first element ~ PCl3 ~ 1 atom Phosphorus + 3 atoms Chlorine NAME ~ MonoPhosphorus trichloride

Classifying Matter 2 large groups:
Pure substance – only one element or compound Mixture – different compounds and/or elements

Mixtures has 2 subgroups:
Homeogeneous: composition always the same no matter what part is observed Heterogeneous: composition is different depending on what part is observed

Understanding of Module 2 Labs
Sand is a covalent compound because it is made of silicon and oxygen which are nonmetals. Corn starch is covalent because it doesn't dissolve in water, and is not a crystal. Baking soda is ionic because it is made up of sodium, a metal, and carbon, a nonmetal. Sugar is ionic because it melts and is soluble in water.

Module 3: Atomic Structure
Electrical charge: Positive vs Negative 1st rule: Like charges repel one another. 2nd rule: Opposite charges attract each other. Every substance has electrical charges. If the # of + = # of -, then it is neutral If more of one charge than the other, the substance takes on that charge

The Nature of Light Since light has no mass and takes up no space, it is not matter. Particle/wave duality theory: light sometimes behaves as a particle and sometimes behaves as a wave.

Electromagnetic Spectrum
The Visual Spectrum When the wavelength is large/wide, the frequency is small/slow. When the wavelength is small/thin, the frequency is large/fast. As a light’s wavelength increases, its energy decreases. As its wavelength decreases, its energy increases.

J.J. THOMSON’S MODEL Cathode-ray tube
He was surprised when he realized that atoms were shooting off something that would bend towards a positively-charged plate! The electron is discovered! “Plum Pudding” Model (electrons stuck at random into a pudding of positive charge

Rutherford’s model Before neutrons were discovered
Created an experiment to test the presence of “alpha particles” Also called the Planetary Model Lacked the concept of light as particles (photons) nor that it moves in various wavelengths

BOHR’S MODEL Bohr developed the idea that electrons were at specific energy levels, and could jump from level to level, but not reside in-between. Like planets orbiting the sun, or runners on a track.

Isotopes Number of Neutrons may vary within the same element
Isotopes behave identically, yet differ in their atomic mass

Our Current Model: Subatomic particles
Proton (positive charge, determines the actual element type) Neutron (neutral charge, different # = different isotope) Electron (negative charge, changing the number of these can create an ion) Relative mass: Proton + Neutron Quarks (protons/neutrons/electrons are made up of these)

Protons (+) – found in atom’s nucleus
Neutrons (0) – found in atom’s nucleus Electrons (-) – found outside atom’s center

Atomic # and Mass # Refer to Worksheet: Atomic Structure Diagrams
Total number of neutrons + protons Total number of protons Refer to Worksheet: Atomic Structure Diagrams Refer to Worksheet: Bohr’s Models Practice

Bohr atom diagrams How do we know this is what carbon looks like? 6 P 6 N

Protons and Electrons equal the number on top of the element in the symbol
Why are there two rings?

The Rings Each row has a different number of rings 1 ring 2 rings

Carbon’s Rings 1 ring 2 rings 3 rings 4 rings 5 rings

QUESTION: Why are there different numbers of electrons in different rings?
Carbon has 2 rings. Why?

ANSWER: Each ring can only hold a maximum amount of electrons as there are elements in that row
1 ring with 2 electrons max in this ring 2 rings with 8 electrons max in this ring 3 rings with 8 electrons max in this ring

Let’s Try Another Element
1 ring 2 rings 3 rings 4 rings 5 rings

Let’s Try Another Element
1 ring with 2 electrons max in this ring 2 rings with 8 electrons max in this ring 3 rings with 8 electrons max in this ring

W/S: Bohr Models Practice
It’s time for: BOHR Models W/S: Bohr Models Practice C Ti P Al

Louis de Broglie’s model
Erwin Schrödinger and Louis de Broglie’s model The question of quantum mechanics: can you know where something is and how fast it is traveling at the same time? Nope. (Uncertainty Principle) We cannot know WHERE electrons are and HOW FAST they are going—we can only know one or the other.

The rest will be in higher levels and orbitals.
Orbital Structure of an Atom Orbitals and the # of electrons 1st orbital - 2nd Orbital - 3rd Orbital - 4th Orbital - 2 electrons 8 electrons 18 electrons 32 electrons The rest will be in higher levels and orbitals.

The math of Atoms 1 1 8 8 92 146 56 26 Fill in the chart below
Two very simple rules: # Electrons = # Protons Protons + Neutrons = Atomic Mass Fill in the chart below Element Electrons Protons Neutrons Mass Hydrogen 1 Oxygen 8 16 Uranium 92 238 Iron 26 30 1 1 8 8 92 146 56 26

Quantum Mechanical Model of the atom
This model is a revision of the Bohr model - There are still orbits in this model and they are called energy levels. - The energy levels are not all Spherical in shape - This model is developed by complex mathematical predictions. The orbitals of the energy levels: S - Orbital - Spherical and holds up to 2 electrons P - Orbital - Dumbbell shapes and holds up to 6 electrons. has three sub orbitals: two electrons for each one

d - Orbital Has several shapes with most notably the dumbbell with donut Holds up to 10 electrons Has five sub orbitals: Two electrons for each one

f - Orbital Has several shapes with most notably the dumbbell with a double donut Holds up to 14 electrons Has 7 sub orbitals: Two electrons for each one

Electron Configuration
Shows which electrons go into each energy level and sub orbital. It is not hard…once the orbit is filled, you fill the next one. There is a specific order in which the energy levels are filled though This is the order they are filled – follow the arrows. The order coincides with the level of each sublevel. It just happens to be that the 4s orbital is less in energy than the 3d orbital. So it will fill first.

It just happens to be that the 4s orbital is less in energy than the 3d orbital. So it will fill first.

Only one electron in the orbital
Let’s try: Let’s start with hydrogen which has 1 electron It would look like this: H – 1s1 Only one electron in the orbital First energy level In the s orbital

Let’s try: This time lets write one for oxygen which has 8 electrons It would look like this: - O – 1s2 2s2 2p4 6. It can hold up to electrons but we only have 4 left 1.First energy level & 1st electron into s orbital 5. Now the p orbital starts 2. In the next s orbital, 2nd electron 3. Next electron goes into the 2nd s orbital, Into the 2nd electron spot 4. In the s orbital, 2nd electron The orbitals fill up like people sitting in a bus: each person sits alone until the benches are full. Then people will double up on each bench.

Orbital Notation O – 1s2 2s2 2p4 1s 2s 2p
Shows which electrons go into each energy level and sub orbital It is more of a picture so that you can see how the electrons are arranged. For example – here is the electron configuration for oxygen again: O – 1s2 2s2 2p4 Orbital notation would go like this: 1s 2s 2p

The orbitals fill up like people sitting in a bus: each person sits alone
until the benches are full. Then people will double up on each bench.

So, Why is there an up arrow and a down arrow?
A longer answer: Several experimental observations can be explained by treating the electron as though it were spinning. The spin can be clockwise or counterclockwise, and so there are two possible values of the 'spin quantum number' that describe the electron. The quantum theory was able to explain the experimental results if the spin quantum number was taken to be either +1/2 or -1/2. An up arrow on an orbital diagram means one of these values, and a down arrow means the other (try not to worry about which is which. It really doesn't matter.) Why do all the parts of the orbitals get one electron each before they add another one? The electrons are negative and have to get the farthest away from each other. Why do each part of the orbital only get two electrons in the 1st place? Well, this one is yet to be explained. For it will be – “because that is just the way it is” !

Practice - Boron (B) - 5 electrons Neon (Ne) - 10 electrons
Magnesium (Mg) electrons Iron (Fe) electrons Chromium (Cr) electrons Challenge Question - Einsteinuim (Es) electrons

One more way to show electron configuration:
We can shorten things up by using a previous element For example: Let’s use Phosphorus with 15 electrons P – 1s2 2s2 2p6 3s2 3p3 Neon 3s2 3p3 P – [Ne]

Module 4: Molecular Structure
Valence electrons: Those furthest from the nucleus highest energy level number it is what effects the chemical behavior of that atom atoms in the same column on periodic table have same # of valence electrons and chemical behavior

LEWIS dot diagrams 1 2 3 4 5 6 7 8 Represents each valence electron
Developed by G. N. Lewis, it forms the basis of current molecule understanding Represents each valence electron 1 2 3 4 5 6 7 8

Lewis made a simpler model for each element
The Lewis Dot Model is concerned with the valence (the electrons in the outer shell) Mg has 2 valence electrons, and would look like this: S has ___ valence electrons, and would look this this: 6

Ne Order of Dot Placement? What is important:
4 8 Ne Other sources: 3 1 7 5 2 6 Apologia Textbook What is important: place one electron on each side before doubling up on electrons

Atoms will combine with other atoms to achieve the octet rule
…making molecules Whoa is me!!! I only have 7 valence electrons! I need 1 more!!!! My hero!! Excuse me, my dear! I’ll help you!! Oh who will help me? Na Cl Cl Na

Ionic Compounds Example: This is Sodium Chloride
Ion: Atom which has gained or lost electrons Loose an electron, which is negative, then the atom becomes an ion and is positively charged Cation: positively charged atom, and keeps same name as original atom Gain an electron, which is negative, then the atom becomes an ion and is negatively charged Anion: negatively charged atom, and “ide” is added Example: This is Sodium Chloride Metals try to loose electrons, while Nonmetals try to gain electrons

Let’s try: Magnesium Fluoride
Positive charges attract Negative charges Let’s try: Magnesium Fluoride Magnesium, with 2 valence electrons, followed the Octet Rule by gaining 6 electrons from 2 different Fluorine atoms: Magnesium Fluoride Its chemical formula is: MgF2

Periodic Table of Ions (See Handout)
All atoms in a given column tend to develop the same charge Exceptions (discussed in a later slide) are found on the periodic table with split boxes Do NOT memorize, learn to find quickly on Periodic Table Usually just remember metals (cations+), nonmetals (anions-)

Ionic Molecole’s Chemical Formula
Steps: Look at periodic table and determine the charge of each element If the charges are the same number, the subscripts for each will be 1 If the charges have different numbers, drop the + and – signs, switch the numbers, and use them as subscripts The chemical formula ratio for ionic compounds must always be the lowest possible ratio To name: use the cation first ________________ Barium Nitride ________________ Aluminum Sulfide

Let’s try: What about this compound?
Beryllium is in group 2A, so it is 2+ Chlorine is in group 7A, so it is 1- Ignore the + and -, and switch the #s What about this compound? Magnesium is in group 2A, so it is 2+ Nitrogen is in group 5A, so it is 3- It’s formula is: It’s name is: Magnesium Nitride

Covalent Compounds Will discuss in Module 5 Will NOT be included
In Module 3 & 4 THT

The Exceptions Cu+1 Cu+2 Cu(I) Cu(II) Transition metals:
Have more than 1 charge Named differently Use Roman Numerals Cu+1 Cu+2 Cu(I) Cu(II)

It’s time for: Lewis Dot Exercises Let’s play

Module 5: Polyatomic Ions & Molecular Geometry
Electronegativity: A measure of how strongly an atom attracts extra electrons to itself Greater desire to gain an electron

THE LARGER THE ATOM’S RADIUS, THE LARGER THE ATOM
ATOMIC RADIUS: THE LARGER THE ATOM’S RADIUS, THE LARGER THE ATOM

How to order elements in a chemical formula?
There's something called Hill System This is the system of writing chemical formulas Carbon atoms are first, then hydrogen atoms Others follow in alphabetical order When the formula contains no carbon or hydrogen atoms, then all the elements are sorted alphabetically.

Homonuclear Diatomic Structures
Homonuclear means the molecule contain only one type of element The most familiar homonuclear molecules are diatomic, meaning they consist of two atoms Not all diatomic molecules are homonuclear Homonuclear diatomic molecules include hydrogen (H2), oxygen (O2), nitrogen (N2) and all of the halogens Ozone (O3) is a common triatomic homonuclear molecule

Covalent Bonds atoms together in covalent compounds
A shared pair of valence electrons that hold atoms together in covalent compounds This is what separates them from ionic bonds which give or take electrons – making ions +/-, with attraction keeping the atoms together

Let’s try: Carbon tetrachloride CCl4
How to order elements in a chemical formula? There's something called Hill System. This is the system of writing chemical formulas. In this system the carbon atoms are first, then hydrogen atoms and then other in alphabetical order. When the formula contains no carbon or hydrogen atoms, then all the elements are sorted alphabetically. Start in the center, and work your way out!!! Center: the atom with the most unpaired electron

Steps to making a Lewis structure:
Count the valence electrons in the entire molecule by adding up the valence electrons/individual atom Put the atom with the most unpaired electrons in the center. If more than one, put all of them in the center, bonded together Place the remaining at0ms around the central atom and bond them with a single covalent bond Fill in the ocetets of the outside atoms. Make sure the center atom has its ocetet.

More Covalent Bonds Cl - Cl O = O N = N
Single Bonds: 1 pair of shared electrons, indicated with a dash (covalent bond) Cl - Cl Double Bonds: 2 pairs of shared electrons, indicated with 2 dashed lines O = O Triple Bonds: 3 pairs of shared electrons, indicated with 3 dashed lines N = N These allow the electron(s) to count for both atoms

Polyatomic Ions -OH -NH4 -NO3 -PO4 -SO4
a charged chemical ion composed of two or more atoms covalently bonded or of a metal complex that can be considered to be acting as a single unit often useful in the context of acid-base chemistry or in the formation of salts.

Let’s try: Calcium Nitrate Ca NO3 - 2+ 2+ Ca NO3 -

VSEPR Theory Lewis Structures only show 2 dimensions
Molecules are 3-demensional (geometric) VSEPR (Valence Shell Electron Pair Repulsion) Theory states: Valence shell electron pair repulsion (VSEPR) theory is a model used in chemistry to predict the geometry of individual molecules from the number of electron pairs surrounding their central atoms. It is also named the Gillespie-Nyholm theory after its two main developers.

EXAMPLES OF VSEPR SHAPES:

Let’s try: PHCl2 1. The basic shape is tetrahedron P Cl H 2. One pair of electrons is not shared with the phosphorus 3. There are now only 3 legs. Cl o Pyramidal with a bond angle of 107 The shape of the molecule is _______________________

Module 6: Changes in Matter & Chemical Reactions
Chemical change: molecules or atoms in a substance are changed Physical change: molecules or atoms in a substance stay the same

Physical Change vs Chemical Change

Physical Phase Changes
Molecules are farthest apart Add more energy, it boils Add energy, it melts Take energy away, it freezes Molecules are closest together

Kinetic Theory of Matter
Matter is composed of a large number of small particles—individual atoms or molecules—that are in constant motion. Motion requires work Work requires energy Energy comes from the surroundings Increasing temperature increases atomic motion Decreasing temperature causes energy to return to surroundings Boiling Point the temperature at which a liquid boils and turns to vapor. Melting Point the temperature at which a solid turns into a liquid.

Density Grams (g) Milliliters (ml)
~the degree of compactness of a substance. Grams (g) Milliliters (ml)

Water is SPECIAL Water occupies MORE volume in its solid phase,
and its density is LESS as a solid than as a liquid. d/t Hydrogen bonding (Module 7)

Balancing Chemical Equations
Chemical Reactions, Counting Atoms, and Balancing Chemical Equations

A Chemical Reaction is…
A process where one or more elements or compounds are changed into one or more different substances A process where the original substance changes into a new substance with new properties But, are new atoms created? Are other atoms destroyed?

Law of Conservation of Mass states that…
Matter is never created or destroyed… Therefore, the mass and types of atoms that make up the reactants must be equal to the mass and types of atoms in the products. The atoms are simply rearranged.

Reactants and Products
Reactant: the original substance (s) Product: the new substance (s) In the chemical equation, A + B AB -the reactants are ‘A’ and ‘B’, which are on the left side of the equation. -the product is ‘AB’ and is found on the right side of the equation. The arrow in the equation means ‘yields’, or ‘produces’.

How Do We Know If It Is A Chemical Reaction?
The following observations provide evidence that a chemical reaction is taking place: The temperature changes A gas is produced (bubbling) A precipitate forms (a solid forms from the combination of two solutions and separates out) A permanent change in color Be aware: Some physical changes produce some of the above observations. For example, food coloring in water and boiling water may look like chemical changes, but are physical changes.

HCl Zn Let’s try: 2 indications that a chemical
Bubbling, and temperature 2 indications that a chemical change has taken place. the atoms are simply rearranged during the reaction

Chemical Equation The arrow means ‘yields’ or ‘produces’.
A chemical equation, or formula equation, represents what happens in a chemical reaction. For instance: The formula equation for the addition of zinc metal and hydrochloric acid looks like this: 2HCl + Zn H2 + ZnCl2 The arrow means ‘yields’ or ‘produces’.

Counting Atoms in Formulas
Before you can correctly write and balance an equation, you must be able to correctly count atoms in chemical formulas. Let’s learn some basics…

Know Your ‘Equation’ Terms
Coefficient – a whole number in front of the formula that represents the relative number of moles of the substance Subscript – the small number to the lower right of an atom that represents the number of atoms of that element in the molecule Coefficient of 2 means there are 2 moles of water 2H2O Subscript of 2 means there are 2 hydrogen atoms in a molecule of water

The subscript ‘2’ belongs to the oxygen only and does not affect the carbon atom. How many atoms are in this formula? CO2 1 mole of CO2 1 atom of carbon 2 atoms of oxygen 3 atoms in this formula

How many atoms are in this formula? 3CH4 3 moles of CH4 1 atom of carbon 4 atoms hydrogen per molecule 15 atoms in this formula

Al(OH)3 How many atoms are in this formula? (OH)3 means there are three molecules of OH bonded to the atom of aluminum. ‘OH’ is in parentheses because the oxygen and hydrogen atoms are a ‘pair’. 1 mole of Al(OH)3 7 atoms in this formula

How many atoms are in this formula? 2Cr2(CO3)3 (CO3)3 means there are three molecules of CO3 bonded to two atoms of chromium Count carefully! 2 moles of 2Cr2(CO3)3 28 atoms in this formula

How To Balance Chemical Equations
Now that we know how to count atoms in formulas, we can begin to balance equations. Let’s look at a balanced equation.

Chemical Equations 2HCl + Zn H2 + ZnCl2
The chemical equation above is a balanced equation. The reactants are: HCl and Zn The products are: H2 and ZnCl2 There are equal numbers and kinds of atoms in the reactants as there are in the products.

H2 H2 O2 Chemical Equations H2O H2O 2H2O 2H2 + O2
A well-written chemical equation must: Represent a real situation – the actual reactants must be able to react and form the products given Contain correct formulas for the reactants and products Abide by the Law of Conservation of Mass, which means that atoms are never created nor destroyed during chemical reactions H2O 2H2O H2 + O2 O2 H2

Chemical Equations Balancing chemical equations is easy if you follow some rules. You can only add a coefficient in front of a chemical formula H2O You cannot change any subscripts, nor add any subscripts O2 You may not place a coefficient in the ‘middle’ of a chemical formula Ca2Cl You must end up with the same number and kinds of atoms on both sides of the equation when you are finished. Count atoms on both sides upon completion to check yourself. 2AB A B

Let’s try: H2 + Cl2 HCl The word equation is:
Hydrogen gas and Chlorine gas yields hydrogen chloride The formula, or chemical equation is: H2 + Cl HCl The above equation does not satisfy the requirements of a balanced equation. Let’s balance it.

Let’s try: Balance the Equation H2 + Cl2 HCl Are the 2 sides balanced?
Begin by counting atoms on both sides of the arrow. Remember that the number and kinds of atoms in the reactants has to equal the numbers and kinds of atoms in the products. It will help if you divide the two sides with a line below the arrow and list the atoms.

Let’s try: Balance the Equation H2 + Cl2 HCl 2 H - 2 Cl - 2 H - 1
You can see that it appears as if a hydrogen and an oxygen atom were lost in this reaction. That violates the Law of Conservation of Mass, so we must balance the equation by adding coefficients where necessary. Let’s add the coefficient ‘2’ in front of the ‘HCl’. Now let’s adjust our atom totals below the equation and check.

Balance This Equation Here’s another example: Al + O2 Al2O3
Al – Al – 2 O – O – 3 Count the number of atoms and list them under the equation. Compare each side. It’s OK if the first coefficient you try does not balance the equation. You might try adding a coefficient of 2 in front of the Al to balance the aluminum atoms. 2Al + O Al2O3 Al – Al – 2 O – O - 3

Balance This Equation Balanced yet??? 2Al + O2 Al2O3 Al – 2 Al – 2
O – O - 3 3 2 How should we fix the number of oxygen atoms? Al – 2 O – 6 Al – 4 O – 6 Add a coefficient of 3 in front of the O2 on the left, and a coefficient of 2 in front of the Al2O3 on the right. Adjust the numbers of atoms below the equation. Balanced yet???

Balance This Equation Oops, we still have a problem.
The number of aluminum atoms do not match. 2Al + 3O Al2O3 Al – Al – 4 O – O - 6 4 Al – 4 O – 6 Remove the 2 in front of the Al, and add a coefficient of 4. Adjust the numbers of atoms. Now we have a balanced equation!

Let’s complete some balancing on our own!!!

Module 7: Describing Chemical Reactions
Now that you know how to balance equations, let’s look at the different types of reactions.

Types of Chemical Reactions
The basic types of chemical reactions are: Synthesis Decomposition Combustion Single Replacement Double Replacement

Synthesis General equation: Example of synthesis: SO2 + H2O H2SO3
Two or more substances combine to form a new compound. General equation: (A and B represent an element, molecule, or compound) Example of synthesis: SO2 + H2O H2SO3

Decomposition Example of Decomposition: 2NaNO3 2NaNO2 + O2
A compound decomposes into two or more simpler substances. Example of Decomposition: 2NaNO NaNO2 + O2

Combustion Example of the combustion of propane:
Combustion: A substance will combine with oxygen, and in the process, give off energy in the form of light and heat. Example of the combustion of propane: C3H O CO H2O

Single Replacement Example of Single Replacement
One element replaces another element in a compound. Example of Single Replacement 3Fe + 4H2O Fe3O4 + 4H2

Double Replacement Example of Double Replacement
Two ionic compounds exchange ions in solution to form two new compounds. Example of Double Replacement FeS + 2HCl H2S + FeCl2

Practice Which of the following are balanced equations? NaCl Na + Cl2
2H2O H2 + O2 2Fe(OH) H2SO Fe2(SO4) H2O 4NH O N H2O

Practice Answers Which of the following are balanced equations?
NaCl Na + Cl2 2H2O H2 + O2 2Fe(OH) H2SO Fe2(SO4) H2O 4NH O N H2O

Practice Balance the following equation: CO2 + H2O C6H12O6 + O2

6CO2 + 6H2O C6H12O6 + 6O2 Practice Answers
Balance the following equation: CO2 + H2O C6 H12O6 + O2 C C O O H H Start by adding a coefficient of 6 in front of the ‘CO2’ and recount the atoms; add a coefficient of 6 in front of the ‘H2O’ and recount; finish by adding a ‘6’ in front of the ‘O2’ and check. 6CO H2O C6H12O O2

Practice Name the type of reaction: 2Na + Cl2 2NaCl K + AgCl Ag + KCl
2HgO Hg + O2 HCl + NaOH NaCl + H2O 2H2 + O H2O

Practice Answers Name the type of reaction: 2Na + Cl2 2NaCl synthesis
K + AgCl Ag + KCl single replacement 2HgO Hg + O2 decomposition HCl + NaOH NaCl + H2O double replacement 2H2 + O H2O combustion

Module 8:Stoichiometry

The Concept of Stoichiometry
The word stoichiometry comes from the Greek words: stoicheion (meaning "element") metron (meaning "measure") Stoichiometry: The study of quantities as it relates to chemical reactions.

Balanced chemical equations can be interpreted many ways…
1N2(g) + 3H2(g)  2NH3(g) N H molecules: H +  N H N H H

1N2(g) + 3H2(g)  2NH3(g) mass: +  28.02g 6.06g 34.08g 34.08g = Mass of reactants and products are always equal!

1N2(g) + 3H2(g)  2NH3(g) moles: +  1 mol N2 3 mol H2 2 mol NH3 This is the relationship we will focus on in this unit!

Mole-Mole Calculations
The coefficients in the balanced equation represent the smallest whole number mole ratio between reactants and products.

Examples of mole ratios for the reaction: 1N2(g) + 3H2(g)  2NH3(g) 1 mol N2 2 mol NH3 3 mol H2 1 mol N2 2 mol NH3 3 mol H2

Example: If 3.86 moles of potassium reacts completely with excess water, how many moles of hydrogen would be produced? hydrogen hydroxide potassium + water  potassium hydroxide + hydrogen __K 2 __ H(OH) 2 __ K(OH) 2 + __ H2 1  + K: 3.86 mol K U: ? mol H2 3.86 mol K 1 1 mol H2 2 mol K 1.93 mol H2 x =

Example: How many moles of aluminum will react with moles of hydrochloric acid? hydrogen chloride aluminum + hydrochloric acid aluminum + hydrogen chloride  2 __Al __HCl 6 __ AlCl3 2 + __ H2 3  + K: mol HCl U: ? mol Al 0.512 mol HCl 1 2 mol Al 6 mol HCl x = 0.171 mol Al

There are three steps needed for these calculations:
Convert gramsmoles (Using the MM of the known.) Convert moles of knownmoles of unknown using the mole ratio Convert moles of unknown grams (Using MM of the unknown.)

Mass-Mass Calculations
Example: What mass of silver chloride will react with 15.0g of aluminum? . __Al 1 __ AgCl 3 __ Ag 3 + __ AlCl3 1  + 1 Al Ag Cl 1 3 1Ag = ) K: 15.0 gAl Review! 3 3 1Cl =35.45(1) U: ? gAgCl 3 143.32g 15.0 gAl 1 1 mol Al 26.98 gAl 3 mol AgCl 1 mol Al gAgCl 1 mol AgCl x x x 239g AgCl =

Mass-Mass Calculations
Example: What mass of oxygen gas would be produced from the decomposition of 37.2g of lithium chlorate? . __Li(ClO3) 2 __ LiCl 2 + __ O2 3  1Li =6.94(1) K: 37.2 gLi(ClO3) 1Cl =35.45(1) 3O =16.00(3) U: ? gO2 90.39g 37.2 gLi(ClO3) 1 1 mol Li(ClO3) 90.39 gLi(ClO3) 3 mol O2 2 mol Li(ClO3) 32.00 gO2 1 mol O2 x x x 19.8 gO2 =

SUMMARY OF IMPORTANT POINTS:
You are calculating the mass for TWO different substances! Molar mass (MM) is ALWAYS equal to ONE mole! The ONLY time the number of moles is different than ONE (1) is when you are doing the MOLE RATIO STEP!

1 mole = 6.02x1023 atoms, molecules, or formula units
Remember… 1 mole = 6.02x1023 atoms, molecules, or formula units 1 mole = 22.4L (for any gas at STP)

More Stoichiometry Conversions
Example: How many grams of aluminum can react with 336L of oxygen? aluminum + oxygen  . Al O-2 aluminum oxide __Al 4 __ O2 3 2  __ Al2O3 + K: 336L O2 U: ? g Al 336L O2 1 1 mol O2 22.4L O2 4 mol Al 3 mol O2 26.98 g Al 1 mol Al x x x 5.40x102 g Al =

More Stoichiometry Conversions
Example: Using the same equation, how many formula units of aluminum oxide can be produced from 137 grams of aluminum? . __Al 4 __ O2 3 __ Al2O3 2  + K: 137g Al U: ? F.U. Al2O3 137g Al 1 1 mol Al 26.98g Al 2 mol Al2O3 4 mol Al 6.02x1023 F.U. 1 mol Al2O3 x x x = 1.53x1024 F.U. Al2O3

Limiting Reactant Calculations
The limiting reactant (L.R.) is the reactant which runs out first and limits the amount of product that can be made. Two mass-mass calculations will be used. The L.R. is the one that yields the smaller amount.

25.4g of potassium fluoride can be produced.
Example: If 17.1g of potassium reacts with 14.3g of fluorine, which reactant is the limiting reactant and what mass of potassium fluoride can theoretically be produced? Word equation: potassium + fluorine  potassium fluoride Formula Equation: 2K + F2  2KF 1K=39.10 1F=19.00 K: 17.1gK 17.1gK 14.3gF2 ?gKF 58.10g U: ?gKF 17.1g K 1 1 mol K 39.10g K 2 mol KF 2 mol K 58.10gKF 1 mol KF x x x = 25.4g KF K: 14.3gF2 U: ?gKF 14.3gF2 1 1 mol F2 38.00gF2 2 mol KF 1 mol F2 58.10gKF 1 mol KF x x x = 43.7g KF Potassium is the L.R. 25.4g of potassium fluoride can be produced.

Other L.R. Calculations Example: How many moles of iron (III) oxide can be produced from the reaction of moles of iron with moles of oxygen? Word equation: Formula Equation: Fe O-2 iron + oxygen  iron (III) oxide 4 Fe + 3 O2  2 Fe2O3 K: molFe U: ? molFe2O3 Can be produced: 13.17 mol Fe 1 2 mol Fe2O3 4 mol Fe x 6.585 mol Fe2O3 = K: molO2 U: ? molFe2O3 18.19 mol O2 1 2 mol Fe2O3 3 mol O2 x = 12.13 mol Fe2O3

Example: Limiting Reactant AND Excess Reactant Calculations
What mass of CoCl3 is formed from the reaction of 3.478x1023 atoms Co with 57.92L of Cl2 gas at STP? How much of the excess reactant reacts and how much is left over? 2Co + 3Cl2  2CoCl3 1Co=58.78 3Cl=106.35 K: 3.478x1023 atomsCo 165.13g U: ? gCoCl3 LR 3.478x1023atomsCo 1 1 mol Co 6.02x1023atomsCo 2 mol CoCl3 2 mol Co 165.13gCoCl3 1 mol CoCl3 x x x = K: LCl2 U: ? gCoCl3 Can be produced: 95.40gCoCl3 ER 57.92L Cl2 1 1 mol Cl2 22.4L Cl2 2 mol CoCl3 3 mol Cl2 x 165.13gCoCl3 1 mol CoCl3 x x = 284.7gCoCl3

What mass of CoCl3 is formed from the reaction of 3.478x1023 atoms Co with 57.92L of Cl2 gas at STP? How much of the excess reactant reacts and how much is left over? 2Co + 3Cl2  2CoCl3 To find out how much excess reactant reacts, do a third calculation using the limiting reactant as the known, and the excess reactant as the unknown. LR ER K: 3.478x1023atomsCo U: ? L Cl2 3.478x1023atomsCo 1 1 mol Co 6.02x1023atomsCo 3 mol Cl2 2 mol Co 22.4L Cl2 1 mol Cl2 x x x = Reacts: 19.41L Cl2

What mass of CoCl3 is formed from the reaction of 3.478x1023 atoms Co with 57.92L of Cl2 gas at STP? How much of the excess reactant reacts and how much is left over? 2Co + 3Cl2  2CoCl3 To find out how much is left over (unreacted), subtract the amount of the excess reactant that reacted from the original amount of excess reactant (from the problem). LR ER 57.92L Cl2 -19.41L Cl2 Left over: 38.51L Cl2

Module 9: Acids & Bases

Acids An acid is an ionic compound that has the H+ ion as its cation.
Acids have a sour taste (Ex: citric acid) Acids will cause indicators to change colors Acids

substances that change color based on ph.
Indicators: substances that change color based on ph.

Acids Turn litmus paper pink or red The pH of acids is less than 7
Acids will react with metals in a single replacement reaction called corrosion. Acids will react with bases in a double replacement reaction called neutralization.

Bases A base is an ionic compound that has the OH- ion as its anion.
Bases have a bitter taste Bases are slippery (Ex: soap) Bases will cause indicators to change colors Bases

Bases Bases change litmus paper blue.
The pH of bases is greater than 7. Bases will react with acids (neutralization)

Review: Naming and Formula Writing for Bases
Use typical rules for naming ionic compounds: NaOH ___________________ Mg(OH)2 ___________________ Aluminum Hydroxide ________ Sodium hydroxide Magnesium hydroxide Al(OH)3 Al (OH)-1

Review: Naming and Formula Writing for Acids
For acids, look at the anion ending! H3PO3 ___________________ HBr ___________________ Carbonic acid ______________ Remember H+1 is always the cation for acids! “phosphite” Phosphorous acid “bromide” Hydrobromic acid “ic” must have been “ate” H2CO3 H (CO3)-2

Recall: Concentration is a measure of the amount of solute dissolved in a given quantity of solvent.

Module 10: Solutions Solute – the substance getting dissolved
Solvent – the substance that is dissolving the solute Acid Solution – mixture of solute and solvent The process of making soda pop: CO2 is the solute

Review:

Ionic Compounds: Metals & Non metals
How Solutes Dissolve Ionic Compounds: Metals & Non metals Water (polar covalent) gets in between the ionic bonds, and pull them away from each other The atoms are going to be floating individually EXAMPLE: NaCl

Ionic Compounds This made the solid change into the aqueous phase
Not ALL ionic solids dissolve in water – the bonds are too great for the water to get between – Insoluble in water Used in antacids, d/t it’s basic properties EXAMPLE: Al(OH)3

Covalent Compounds: nonmetals
Pure Covalent Solids cannot dissolve in water as they have no electrical charges Polar Covalent Solids can dissolve in water into individual molecules that made it up, NOT individual atoms The water molecules surround the individual covalent molecules, and the solid becomes aqueous EXAMPLE: C2H4O2

Concentration Indicates the amount of solute dissolved in a given quantity of solvent. Dilute: a small amount of solute Concentrated: a large amount of solute

Solutions Recall, aqueous is a solution where water is the solvent.
Example: NaCl(aq) Na+ + Cl- Solubility: The maximum amount of solute dissolved in a particular solvent at a specific temperature. Saturated: No more solute can dissolve Unsaturated: More solute can dissolve Supersaturated: More solute is dissolved than theoretically possible

Factors that affect solubility:
Temperature: most solid substances have higher solubility as temperature increases All gas solutes have lower solubility as temperature increases Pressure: only affects gas solutes All gas solutes have higher solubility as pressure increases To make a supersaturated solution: Add more solute than solubility allows Heat the solution up Slowly cool it down This is a temporary and unstable state for a solution!

Rate can be increased by:
Dissolving Rate of dissolving: How fast a solute dissolves in a solvent – not to be confused with how much. Rate can be increased by: Increasing temperature- There is more kinetic energy available to meet the activation energy (energy available for dissolving) Stirring- Increases the interaction between solute and solvent Powdering- Increase surface area of the solute which increases the interaction between solute and solvent

Liquid-Liquid Solutions
Miscible: two liquids which uniformly mix together (ex: milk and water) Immiscible: two liquids which will not mix, forms two layers (ex: oil and water) Non-polar + non-polar = miscible Polar + Polar = miscible Non-polar + Polar = immiscible As a general rule: “Like dissolves like”

Steps for making a 0.5M solution
Add 0.5mol of solute to a 1.0L volumetric flask half-filled with distilled water

Steps for making a 0.5M solution
2. Swirl the flask to dissolve the solute. 3. Fill the flask to the mark etched on the side of the flask (1-L mark).

Neutralization Reactions
A neutralization reaction is a double replacement reaction between an acid and a base to produce water and a salt (an ionic compound).

Examples Predict the products in words and formulas, and balance! Predict in formulas only and balance. __ Ca(OH)2 + __ H3(PO4)  hydrogen sulfite sulfurous acid potassium hydroxide potassium sulfite +  water + H(OH) Ca+2 (PO4)-3 3 2 __ H2O 6 + __ Ca3(PO4)2 1 3 1231 Ca 3112 6 (OH) 6 6 H 6 2 (PO4)

Molarity Important: The volume is the total volume of resulting solution, not the solvent alone. Molarity= Moles of solute Liters of solution Important: The volume is the total volume of resulting solution, not the solvent alone.

Performing Calculations With Molarity
Example: What is the molarity of a solution that contains 0.25 moles of NaCl in 0.75L of solution? M= mol= L= ? 0.25 mol NaCl 0.75L 0.25 mol NaCl M = 0.75L Molarity = 0.33 mol/L or M

Example: What volume of a 1.08M KI solution would contain moles of KI? M= mol= L= 1.08 M 0.642 mol KI ? 0.642 mol KI 1.08 M = L Volume = 0.594L

Example: How many grams of CaBr2 are dissolved in 0.455L of a 0.39M CaBr2 solution? M= mol= CaBr2 L= 1 Ca=40.08 0.39 M 2 Br=79.90(2) mol 199.88g 0.39 M ? = 0.455L 0.455L mol = 0.18 mol CaBr2 199.88gCaBr2 0.18 mol CaBr2 x = 36g CaBr2 1 mol CaBr2

Titrations A known quantity of one solution is measured out, and the other solution is added from the buret until the two solutions have neutralized each other. The point at which two solutions have neutralized each other is called the endpoint of the titration. An indicator is used to mark the endpoint of the titration. Phenolphthalein: Is an indicator often used in acid-base titrations. It is colorless in acid and pink in base.

Titration Calculations
To calculate the unknown concentration of an acid or base using titrations, the following steps are used: Write the balanced chemical equation. Convert the known molarity into moles using the molarity formula: M=mol/L. Convert moles of known into moles of unknown using the mole ratio. Using the molarity equation, use moles of unknown and divide by volume to get molarity. Titration Setup

Example 0.025L of a 0.28M HF solution is neutralized by 0.047L of an unknown Mg(OH)2 solution. What is the concentration of the Mg(OH)2? ___ HF + ___ Mg(OH)2  H(OH) 2 1 2 ___ H2O + ___ MgF2 1 0.025L 0.047L 0.28M ? M Step 1: Step 2: mol 0.007mol HF 1 mol Mg(OH)2 0.28M = mol Mg(OH)2 x = 0.025L 2 mol HF mol HF = 0.007mol Step 3: mol = 0.074M Mg(OH)2 0.047L

Do not need to memorize…need later in physics
Module 11: The Gas Phase pressure: the force per unit exerted on an object. Pressure equation: Do not need to memorize…need later in physics

Gas molecules: The unit for force is the NEWTON The unit for area is m
2 The unit for area is m 2 Pressure is Newtons/m This is the Pascal (Pa) Gas molecules: move around quickly collide with anything that gets in their way when they collide, they cause a force on that object the average force per unit of area is called the gas pressure

Identified in textbook

Ideal Gas Law

Ideal Gases: compared to the total volume available to the gas.
The molecules/atoms that make up the gas are very small compared to the total volume available to the gas. The molecules/atoms that make up the gas must be so far apart from one another that there is no attraction or repulsion between them. The collision that occur between the gas molecules/atoms must be elastic. No energy can be lost in the collision. No energy can be lost when colliding with the walls of the container.

These are used as a reference to determine if a gas is ideal or not.

When 2 or more gases are mixed together, the total pressure of the
mixture is equal to the sum of the pressures of each individual gas.

Vapor Pressure

Boiling Point: ~the temperature at which the vapor
pressure of a liquid is equal to normal atmospheric pressure.

Module 12: Energy, Heat, & Temperature
Energy: the ability to do work. Work: the force applied to an object times the distance that the object travels parallel to the force Heat: energy that is transferred as a consequence of temperature differences.

1st Law of Thermodynamics
Energy can NOT be created or destroyed. Energy can ONLY change form. God created only so much energy for our universe. Potential Energy Kinetic Energy

Measuring Energy Energy is measured in Joules: = 1 gram = 102 grams
102 paperclips Lift 102 paperclips straight up at a constant speed for 1 meter = 1.0 Joule

Measuring Heat Heat is measured in Calories: 1 gram water
One degree Celcius How much Heat is needed to:

Relationship between Energy & Heat
Joules and Calories basically measure the same thing 1 calorie = Joules Food calories do NOT equal chemistry calories Equals 1,000 chemistry calories Food Calories

Like density, each substance has a unique specific heat.

Specific Heat One specific heat needs to be memorized: The specific heat of water is cal/g C or J/g C Unit of heat: Joules Unit of mass: gram Unit of temperature: Celsius

Calorimetry -qobject = qwater + qcalorimeter Measured by a calorimeter
Calorimetry equation -qobject = qwater + qcalorimeter q = amount of heat gained or released - in front of q = the object was releasing heat, loosing energy

Module 13: Thermodynamics
This is the branch of science that deals with the relations between heat and other forms of energy (such as mechanical, electrical, or chemical energy), and by extension, of the relationships between all forms of energy. Watch this video on Thermodynamics, and fill out the worksheet on igrade

1st Law of Thermodynamics
Potential energy – energy that is stored Kinetic energy – energy in motion Enthalpy (H) – the energy stored in a substance H – energy change that accompanies a chemical rx.

enthalpy 1 calorie = 4.184 Joules Most chemical rxs are in kcal
H is negative H is positive Heat is released Heat is absorbed HCl + NaOH H2O + NaCl + energy 2NH4SCN + Ba(OH)2(H2O)8 + energy 2NH4OH + Ba(SCN)2 + 8H2O Energy is a product in an exothermic rx Energy is a reactant in an endothermic rx

Let’s Try: CH4 + O2 CO2 + H2O CH4 + 2O2 CO2 + 2H2O + 803.1 kJ
When CH4 undergoes complete combustion, the change in enthalpy of the reaction is kJ. Write a balanced chemical equation representing this process. Include energy in your equation. HINT: Combustion refers to reacting a substance with O2 to produce CO2 and H2O CH4 + O2 CO2 + H2O CH4 + 2O CO2 + 2H2O kJ Unbalanced chemical equation H is negative, so Rx is exothermic!

enthalpy Exp. 13.1 Can You Determine the H of a Chemical Rx?
Equation used: q = mc T q is the amt of heat transferred m is the mass of the sample being heated c is the specific heat capacity of thing being heated T is the change in temperature Specific heat of vinegar is 41 J/g C Vinegar added to calorimeter Lye added to calorimeter Temperature increases Complete the formula for amt of heat transferred Calorimeter (2 nested cups)

Determining H Using Bond Energies
In a chemical reaction: The valence electrons of the reactants must rearrange themselves into configurations necessary to make the products Requires a lot of chemical bond breaking & chemical bond making Lewis Structures: EXAMPLE: CH O2 CO2 + 2H2O All bonds on reactant side - broken All bonds on product side - remade Energy is absorbed Energy is released

In these chemical reactions, energy is being absorbed & released Bond energy – the strength of a bond - unit is kJ/mole It is about: how much energy is required to break 1 mole of those bonds -OR- how much energy is released when 1 mole of these bonds is formed The only way to know which bonds are breaking and which are forming is by looking at the Lewis Structures

EXAMPLE: 1 mole of Cl2 The Lewis Structure: There is only 1 single bond According to the table, the energy of that bond is 239 kJ/mole To break Cl2 apart involves just 1 mole – 239 kJ/mole To form Cl2 involves 2 moles of Cl coming together – 239 kJ/mole Do NOT need to memorize table – just know where to find it!

H = Energy required to break bonds – Energy released when bonds form The only way to know which bonds are breaking and which are forming is by looking at the Lewis Structures If the reactants absorbed more energy (required more energy) than was released by the products – ENDOTHERMIC (result is positive) If the reactants absorbed less energy than was released by the products – EXOTHERMIC (result is negative)

Hess’s Law EXAMPLE: 2 different ways of making water 2H2 + O2 2H2O
Famous Russian chemist and physician Enthalpy is a state function - any quantity that depends solely on the final destination, not on the way we got there The amount of energy contained in a substance is independent of how it is made EXAMPLE: 2 different ways of making water 2H2 + O2 2H2O 2H2O H2O + O2 Enthalpy of formation ( ) - a reaction that forms a single substance from simpler components The Enthalpy is equal, using 2 different paths

Hess’s Law Having the of a substance is equivalent to having the enthalpy of that substance Standard Conditions - room temp (25 C) and 1.00 atm Standard Conditions STP Some are exothermic and some are endothermic depends on the state of the substance (vapor vs. liquid) Ca (s), Na (s), K (s), Fe (s), H2 (g), N2 (g), O2 (g), and Cl2 (g) = O Use Hess’s law when you have the of each substance in the equation Use Bond Energies whenever you do NOT have every necessary = summation

High school science course 1 credit/unit with lab

Similar presentations

Presentation on theme: "High school science course 1 credit/unit with lab"— Presentation transcript:

Similar presentations

About project

Feedback

Log in

Auth with social network:

High school science course 1 credit/unit with lab

Similar presentations

Presentation on theme: "High school science course 1 credit/unit with lab"— Presentation transcript:

Similar presentations

About project

Feedback