Chemistry Midterm

Chemistry

• Testing

• Hypothesis

• Conclusions

• Methods

• Laws

• trial and error

• periodic table

• Elements

• atomic numbers and weights

• Mol

• Noble gasses 

• Halogens

• Alkalines

• Outer shells

• Protons

• Neutrons

• Electrons

• Covalent bonds

30 August 2024

What types of things do you measure in your daily life?

• Temperature

• Ingredients 

• Food

• Time

• Speed

• Qualitative: hunger, thirst, emotions

How do you measure those things?

• Thermometer

• Measuring cups

• Speedometer 

• Odometer

• Teaspoons

• Tablespoons

What types of things might we measure in chemistry this year?

• Weight

• Heat

• Size

• Width

• Time

• Chemicals

• Volume

• Energy

• Mass

• Acidity

How might measuring those things be different?

• Have to be more precise

• What tools we use

• Different measuring systems, grams, kilograms, etc. 

Scientific Method

• Research

• Hypothesis

• Process

• Conclusion

• Observations

• Reasoning

• Experimenting 

• Steps

• Instructions

• Order

• Educated guess as to what your answer to your question is

• Idea

• Don’t use words like proves, or disproves—shows, demonstrates

• Be flexible depending on what happens

• Order does not have to be exact

1. Make an observation

2. Ask a question

3. Do background research

4. Form a hypothesis

5. Test your hypothesis with an experiment

   1. Variables:

      1. Independent variable: Part of the experiment that is changed by the experimenter; it affects another part (the thing that you change/are testing)

      2. Dependent variable: Part of the experiment that changes due to the independent variable (what is measured/observed as a result of the change)

      3. Standardized variables: Variables that are not changed (each test and experiment this variable stays the same)

6. Analyze your data and draw a conclusion

7. None of this matters unless you communicate your results and share your new knowledge of the world

What is important while making a measurement?

1. Precision

2. Accuracy

• All measurements contain some amount of error.

• When scientists make measurements, we evaluate both the accuracy and precision of the measurements

• Accuracy refers to how close a measurement is to the expected value

• Precision refers to how close a series of measurements are to each other

• Error is defined as the difference between an experimental value and an accepted value

• Percent error expresses the error as a percentage of the accepted value

• % error=[experimental-accepted]100/accepted 

• Avg=34.58      %error=[34.58ml-34.1ml]100/34.1ml =1.4%

• His measurements are all close together, which demonstrates the accuracy of the experiment

Significant Figures

• The precision of a measurement is indicated by the number of digits

• Reported digits are called significant figures and include all known digits plus one estimated digit

Recognizing Significant Figures

1. Nonzero numbers are always significant 

2. Zeros between significant figures figures are always significant 102, 100002

3. All final zeros to the right of the decimal are significant 1.0, 1.00

4. Placeholder zeros are not significant 160ml, 0 is not a sig fig between 100-200

1. 23.3

2. 1.37

3. Accurate, not very precise

4. Not accurate, very precise

5. Data set 2 is very accurate, and data set 3 is very precise

202=3

1.00=3

150,000=2

0.000001230=4

Ex. round 2.532 to 3 sig figs=2.53

Ex. round 2.536 to 3 sig figs=2.54

Multiply and divide

4.320 cm times 1.6 cm=6.912=6.9cm^2

Scientific notation

• Everything before the times 10 is significant

• Number larger than one, the exponent will be positive

• Large number-between 1-10

• 2.0289 * 10^4

• 1.000002 * 10^6

• 1250000000000000=1.25 * 10^19

• Small number-between 0-1

• 0.00527 =5.27 * 10^-3

• 0.0000000082=8.2e-9

• 3*10^6=3000000

• 6.26*10^9=6260000000

• 5*10^-4=.0005

• 8.45*10^-7=.000000845

• I would measure the coffee mug with a ruler for the radius to then find out the circumference, and measure the height of it. I would use a scale to measure the weight. And measure how much liquid could fit in it without overflowing. 

Measurement types

SI Units

• To better describe the range of possible measurements, we add prefixes to the base units

• These prefixes are based on factors of 10 and can be used on all SI units

• To convert between different prefixes, we will multiply or divide by factors of 10 (move the decimal place to the right or left)

• atto=10¹⁸

• femto

• pico 

• μ=micro (placeholder for base unit)

• U=unit youre using  (liters, seconds, cm)

• deka=10 times bigger than a meter

• hecto=100

• kilo=1000

• mega=10000

• giga

• tera

• peta

• exa

• M…k.h.da.U.d.c.m…μ

• King henry doesnt usually drink chocolatte milk

• U is your base unit (m,L,g,s)

• Move the decimal once between kilo and milli

• Move the decimal three times between mega (M) and kilo (k) and between milli (m) and micro (μ)

1. 2.47 km to dm=24700 dm

2. 96.35 μL to L=.00009625 L

3. 500 cg to Mg=.000005 Mg

4. 1.23 m (base unit meter-right) to mm=1230 mm

Command period-exponent

Command comma-subscript

Cubic centimeter-Millilter 

Scientists count atoms and molecules in groups: Avogadro’s number

1 mole=6.02 x 10²³ atoms or molecules

• In chemical reactions, what matters is the number of atoms of molecules of a substance present, NOT the mass.

• Moles of different substances will have different masses

• A mole is a group just like a dozen or a gross, it is just A LOT BIGGER.

How big is a mole?

• 1 mole of marbles is enough to cover the entire earth to a depth of five miles

What is a mole? Why do scientists use moles to count atoms?

• A mole is a way to group atoms together to measure them with more accuracy

• Scientists use moles because they are an easier way to count groups of atoms instead of individual atoms, which would be a lot of counting. 

Dimensional analysis

• Is a method of converting units by using conversion factors

1. inches -feet=12in=1f 12in/1ft 1ft/12in

2. seconds-minutes=60s=1min 60s/1min 1min/60s

3. A speed of 25 m/s= 25m=1s 25m/1s 1s/25m

4. Millimeters to meters=1000mm/1m 1000mm/1m 1m/1000mm

• Multiplying a quantity by a conversion factor changes the units of the quantity, but does not change its value

• Conversion factors must cancel one unit and introduce a new one

• More than 1 conversion factor may be required

=72 eggs

8.50 s to min:

` =0.1416666667 min=0.142 min (3 sig figs cause 8.50)

14 ft to cm  (1 inch=2.54 cm):

=430 cm

6.75 g AL to atoms AL (1 mol=6.02 x 10²³ atoms)

=(6.75)(6.02e23)/26.982 =1.506004003e²³=1.51 x 10²³ atoms of Al

How many seconds are in 12 days?

=1036800s= 1.0 x 10⁶

What is the mass of 3.50 x 10²³ magnesium atoms?

=14.1 g/Mg

What is Matter?

• Matter is made of atoms

• Matter comes in different states (ex. solid, liquid, gas)

• Matter takes up space with volume

• Matter has mass

States of Matter: 

Solids

• Matter that has a definite shape (cannot flow) and a definite volume

• Particles are tightly packed

• Not compressable

Liquids

• Matter that has a definite volume, but takes the shape of its container (flows)

• Particles are less tightly packed

• Not really compressable

Gasses

• Matter that takes the shape of its container and fills the entire volume of its container

• Particles are far apart

• Very compressable

Pure Substances 

• A pure substance has a uniform and unchanging composition (ex. water, gold bar, dry ice)

• Mixture: multiple types of particles

Atoms, Elements, and Compounds

• Atom: smallest unit of an element that still maintains the chemical identity of the element

• Element: pure substance that cannot be broken down into simpler, stable substances; made of one type of atom

• Compound: pure substance made by two or more atoms of different elements joined by chemcal bonds (ex. H₂o)

Mixture

• A mixture is a combination of two or more pure substances in which each substance retains its individual chemical properties

• Most everyday substances occur as mixtures, and it is difficult to keep any substance pure

Types of Mixtures

•  A heterogeneous mixture is a mixture that does not blend smoothly throughout, and in which the individual substances remain distinct (ex. taco, italian dressing, mnms)

• A homogeneous mixture or solution is a mixture that has constant composition throughout; it always has a constant phase (ex. coffee, air, ocean, galvanized square steel, brass)

Physical Properties of Matter

• A physical property is a characteristic that can be observed or measured without changing the sample’s composition

  • Density, color, odor, hardness, melting point, boiling point

• Extensive properties depend on the amount of substance present

  • Mass, length, volume

• Intensive properties do not depend on the amount of substance present

  • Color, odor, hardness, density

Density is an Intensive Physical property

• As the amount increases, both mass and volume increase, but the ratio (the density) stays the same

Chemical properties of matter

• The ability of a substance to combine with or change into another substance is called a chemical property

Physical changes

• A physical change alters a substance without changing its chemical composition

Physical changes-Phase changes

• One type of physical change is a phase change, which is a transition from one state of matter to another

Phase Changes

• Solid-Liquid

• Melting=Solid to Liquid

• Freezing=Liquid to Solid

Phase Changes

• Liquid-Gas

• Evaporation=Liquid to Gas

• Condensation=Gas to Liquid

Phase Changes

• Gas-Solid

• Sublimation=Solid to Gas

• Deposition=Gas to Solid

Chemical Changes

• A process that involves one or more substances changing into new substances is called a chemical change (or a chemical reaction)

• A chemical change ALWAYS results in a change in properties

What are the differences between physical and chemical properties?

• A physical property is a characteristic that can be observed or measured without changing the sample’s composition

  • Density, color, odor, hardness, melting point, boiling point

• The ability of a substance to combine with or change into another substance is called a chemical property

What are the differences between intensive and extensive properties?

• Intensive properties do not depend on the amount of substance present

  • Color, odor, hardness, density

• Extensive properties depend on the amount of substance present

  • Mass, length, volume

What are the differences between physical and chemical changes?

• A physical change alters a substance without changing its chemical composition

Physical changes-Phase changes

• One type of physical change is a phase change, which is a transition from one state of matter to another

Chemical Changes

• A process that involves one or more substances changing into new substances is called a chemical change (or a chemical reaction)

• A chemical change ALWAYS results in a change in properties

  • Fire, color change, gas produced/bubbling, 2 liquids form a solid=precipitant, temperature change without touching it

• BONUS!!!!!!!!

  • David Baker-computational protein design

  • Demis Hassabis+John M. Jumper-protein structure prediction

  • Super sexist and racist Nobel Prize-big gender gap

• Phases of Matter

  • The states in which matter exists (solid, liquid, gas)

  • Any part of a system that has uniform properties and competition

  • In a solid there is not a lot of movement in the molecules, remain the same shape

  • In a liquid, particles not so tightly packed together, can flow past each other, therefore, take shape of their container,

  • In a gas, particles move far apart very quickly, take shape and volume of their container 

• Types of vaporization

  • Liquid to gas=vaporization

  • Evaporation-Heat (energy) in the air adds kinetic energy to the molecules at the top of the water making them move faster. They break the attractive forces of the water and become a gas

  • Boiling-Directly changing the conditions. Heat is added directly. Gas forms within the liquid and then escapes. 

• Heat of Fusion and Heat of Vaporization 

  • Heat of Vaporization (ΔH_vap°)

    •  Amount of heat needed to turn 1 gram of a liquid to a vapor.

  • Heat of Fusion (ΔH_fus°)

    • Amount of heat needed to convert 1 gram of solid to a liquid. 

• Phase Diagrams 

  • Phase Diagram-A graph of pressure vs temperature that shows the conditions under which the phases of a substance exists.

  • Triple point-the temperature and pressure conditions at which the solid, liquid, and gas of the substance can coexist at equilibrium

    • Where all things can be true, cold enough to freeze, warm enough to solidify

  • Critical point-shows the temperature and pressure where a substance can no longer exist as a liquid. 

    • Too hot to be a liquid, but too dense to be a gas

Critical Point: Too much pressure to be a solid and too hot to to be a liquid—-creates a supercritical fluid

Specific Heat Capacity

• The more heat an object gains, the hotter it gets

  • q (heat(J))=c (specific heat capacity, a constant for every substance(J per gram degrees C)) x m(mass(g)) x Delta(change in (degrees C, temperature (t_f-t_j)

  • Example: How much heat is needed to warm 250 grams of water from 22degreesC to 98degreesC? 

  • ROUND TO LEAST SIG FIG

    • q=(4.184J/gxc)(250g)(98degreesC-22degreesC)

    • q=(4.184J)(250)(76)

    • q=79000 J

• The specific heat capacity, c, is the amount of heat required to raise 1 g of a substance 1K or 1 degree C

Specific Heat

• The heat required to raise the temperature of the 1 gram of a specific substance by one degree.

  • High Specific Heat=

    • Heat goes in and out SLOWLY

  • Low Specific Heat=

    • Heat moves in and out QUICKLY

  • Metals have low specific heat capacity ex. that's why we use them in our ovens

Solid, liquid, and gas all have different specific heat capacities

Pure Substances: Elements and Compounds

• An element is a pure substance that cannot be seperated into simpler substances by physical or chemical means.

• Each element has a unique chemical name and symbol

Compounds

• A compound is a pure substance made up of two or more different elements that are combined chemically

• Properties of compounds are different than the properties of the individual atoms

Seperating Compoounds into Elements

• Compounds can be separated into elements by chemical processes

• Seperating compounds often requires external energy, like heat or electricity

• In general, compounds that occur in nature are more stable than individual substances

Seperating Mixtures

• The substances in a mixture are physically combined

• The processes used to separate a mixture are physical processes based on 

Filtration

• Heterogeneous mixtures composed of solid+liquid can be separated by filtration

Distillation

• Homogeneous mixtures composed of liquid+liquid can be separated by distillation, based on the differences in boiling point

Crystalization

• Homogeneous mixtures composed of solid+liquid (solid dissolved in liquid) can be separated by crystallization

• When a solution holds as much dissolved substance as it can, the addition of even a tiny bit more causes the solid to come out of the solution

Chromatography

•  Chromatography separates components of a mixture based on the ability of each component to travel across the surface of another material

Kenetic Molecular Theory

• All matter is made out of particles that are in constant motion

• Kenetic Energy is…

  • The energy of motion

  • KE=½ mv²

The Intermolecular Force is…

• The attraction of one molecule to another molecule, not bonds

Solids and Kinetic Molecular Theory

• Flow:

  • Does not flow-particles are too close together 

• Particle motion and Kenetic Energy

  • Vibrational energy, low kinetic energy energy, strong intermolecular forces (cold-vibrate a little bit, hot-vibrate more and break molecular forces to slide around eachother)

Liquids and Kinetic Molecular Theory

• Flow: 

  • Flow: Particles move past each other

• Particle motion and kinetic energy

  • Particles interact often with more kinetic energy than a solid, but there are still strong intermolecular forces

• Role of attractive forces

  • The forces hold the particles together allowing for flow

Gasses and Kinetic Molecular Theory

• Particle motion and kinetic energy

  • Particles can move past each other in order to flow (lots of space in between particles)

Assumptions about gasses in Kinetic Molecular Theory

• Size and volume

  • Gas particles are hard round spheres that do not take up space

  • Gas particles are so tiny, particles essentially do not take up any space. Take up .1% of volume of container they are in

• Attraction to each other

  • Gas particles are not attracted to each other (no intermolecular forces)

• Motion

  • Gas particles move in constant random motion

• Collisions

  • Gas particle collisions are perfectly elastic (no kinetic energy is lost in the system)

Lucy: 26.4

After 30 sec: 29.3

Maya: 27.4

After 30 sec: 29.4

i mix my beads like i mix my beverages

• Particles were moving faster, therefore made a greater sound and hit the sides of the cup more, collide more-hit harder

Temperature

• Temperature is determined by…

  • A substance’s average kinetic energy for all the molecules

• Average Kinetic Energy

  • The kinetic energies of all the different particles are averaged together

  • Temperature/kinetic energy/motion relationship:

    • Higher temperature materials will have a wider range of energies

Temperature (2)

• Kinetic energy relationship to Kelvin:

  • They are directly proportional

• Absolute zero value is…

  • -273 degrees or 0K and it means

    • All molecular movement stops

Solid Structures

• A crystal is

  • A material whose particles are arranged in an orderly, repeated 3-dementional pattern

• A unit cell is

  • The smallest group of particles within a crystal that retains the geometric shape of the crystal

• A crystal lattice is a 

  • Repeating array of unit cells

Solid Structures (2)

• Allotropes are 

  • Two or more different molecular forms of the same element in the same physical state

• Amorphous solids are

  • A solid without a crystal form; it lacks an ordered internal structure

Gas behavior

• Diffusion is

  • The tendency of molecules from a higher concentration to a lower concentration until the concentrations are equal

• Effusion is

  • Gas moves through tiny spaces into areas of lower pressure

  • Small particles are better at this, easier for tiny things to go through tiny things

Gas Behavior

• Gas pressure

  • The sum of the forces exerted on the surface area of an object

  • It is caused by the many different air molecules colliding with the surface

• A vacuum is…

  • An empty space with no particles or pressure

• Atmospheric pressure is…

  • The collision of atoms and molecules in the air with the surface of an object

Heterogeneous Aqueous Solutions

• A suspension is

  • A mixture where the particles of one material are much larger than the particles of the other material

  • It often includes particles in the solid and liquid phases

  • An example is…sand and water

• Homogeneous solution (water), particles too small to have light bounce off of it, cant see light in it

• Heterogeneous solution bc some of the particles are bigger (jell-o), particles are thicker and therefore the light can bounce off of the thicker particles

Heterogeneous aqueous solution 

• A colloid is

  • A heterogeneous mixture made of large particles that are spread throughout

• It includes particles in the solid, liquid, and/or gas phases

• An example is glue, fog, jell-o

Heterogeneous aqueous solution 

• The tyndall effect is…

  • The scattering of visible light by particles in a colloid

• Brownian motion is…

  • The chaotic movement of colloidal particles caused by the collision of molecules

Heterogeneous aqueous solution 

• Coagulation is…

  • The clumping of charged colloidal particles with other charged particles

• An emulsion is…

  • A colloid in which polar and nonpolar molecules are brought together with the help of an additional particle (emulsifying agent)

• An example is…mayonaise, whipped cream

Video Notes

• 440 BCE Democritus first said that everything is made up of little particles surrounded by empty space, they vary in size and shape depending on substance they compose

• Atomos is the greek word for indivisible 

• Ideas were first opposed by people like Aristotle

• Aristotle said that matter was made of four elements, Earth, Wind, Fire, and Water

• 1808, a quaker teacher named John Dalton wanted to challenge Aristotle’s theory

• Dalton showed that common substances always broke down into the same elements in the same proportions

• His conclusion was that the different compounds were composed of different elements atoms of different mass and size based on first law of thermodynamics

• Atomic theory

• JJ Thompsons discovery in 1897 was the electron

• He showed atoms uniformly packed spheres of positive matter filled with negatively charged electrons

• Nobel prize in 1906

• Ernest Rutherford “Father of Nuclear age”

• He shot small, positevely charged, alpha particles at a sheet of gold foil

• Under Thomson’s model, the positive charge was spread evenly so it was not enough to deflect the particles in any one place, it would’ve gone right through, but some bounced back, foil was like a net

• Concluded that atoms consisted largely of empty space with a few electrons, and most mass is in center, he called it the “nucleus”

• The alpha particles passed throught the gaps but bounced pack from the dense, positively charged nucleus

• 1913, Niels Bohr said that electrons orbit the nucleus at “fixed energies and distances”

• They are able to jump from one level to another, but cannot exist in the space between

• Electrons simultaneously behaved as waves, they do not have an exact location

• “Uncertainty principle” Werner Heisenberg demonstrated that its impossible to determine the exact position and speed of electrons at any given time as they orbit an atom

• From the range of possibility grew the quantum model of an atom

• Firework: electrons circle atom shift between energy levels, they absorb/release energy in wavelengths of light

What are Quarks?

In 1964 Murray Gell-Man and George Zweig theorized that particles could be described by blends of three particles. Fundamental means that the particle's bonds cannot be broken. Quarks need to have a fractional charge, they make up all particles. An experiment was conducted where electrons were fired at protons and the scientists observed that the electrons bounced off protons because of the particles inside them. Those added to the evidence that quarks exist. At the time that the experiment was conducted, they had "up quarks" which had a charge of + 2/3, and "down quarks" with a charge of - 1/3. They have different masses, charges, and spin. They are the most common ones, and smallest. Protons are made of 2 up quarks and 1 down quark. Protons=+1. The math behind it is 2(up quarks, + 2/3 +2/3= 4/3-1/3=3/3). A neutron is 2 down quarks and 2 up quarks, which cancel each other out to become a neutral charged particle. There are different types of quarks beside up and down ones, such as, strange quarks, charm quarks, top quarks, and bottom quarks. Up and down quarks were named that because of their spin. Strange quarks were named that because they were present in particle decays that lasted longer than their estimated life span. Charm quarks are called charm because it intrigued the physicists studying them. Bottom and Top quarks were named by a physicist named Harari, and chose them because they go with the up and down quark names. They exist together, and people can observe them because of particle accelerators.

The Atom

• Mass of an atom comes from the nucleus

• 118 different types of atoms=118 elements

What makes an atom of one element different from an atom of a different element?

• Atomic number is the number of protons in an element

• In a neutral atom (total charge is 0), the number of protons=number of electrons

• All elements are defined by the number of protons they have

• Recall that protons and neutrons both have a mass of 1 amu (atomic mass unit)

• The atomic mass (atomic weight) = number of protons + number of neutrons

• All atoms of a single element have the same number of protons, but the number of neutrons can vary

• Atoms with the same number of protons and different numbers of neutrons are called isotopes

• The atomic  mass listed on the periodic table is an average of all of the possible isotopes 

Hyphen Notation

• Sodium

  • Atomic Number=11

  • Atomic Mass=22.99

• Write the name of the element, a hyphen and then the atomic mass

  • Sodium-23

• If you have a sodium atom with 11 neutrons

Nuclear Symbol Notation

• Fluorine

  • Atomic Number=9

  • Atomic mass=19.00

• Write the atomic mass over the atomic number and then the symbol

Ions

• Neutral atoms can become charged particles called ions by gaining or losing electrons

• If an atom loses one or more electrons, it becomes positively charged ion called a cation

• If an atom gains one or more electrons, it becomes a negatively charged ion called an anion

Gram to mole conversions

• What is a mole? How many particles in a mole?

1 mole = 6.02 x 10²³ atoms or molecules

• There is exactly one mole of atoms in the atomic mass of an element when that mass is expressed in grams.

Setting up Gram to Mole Conversions

1. Step 1 – Determine what you are given and what you are looking for, or the unknown

2. Step 2 – Set up a dimensional analysis problem using molar mass as your conversion factor

3. Step 3 – Solve the equation & evaluate your answer

What is the mass in grams of 3.50 mol of the element copper, Cu?

Step 1:  Given ___3.50 mol Cu_____________      Unknown g Cu_______

Step 2:   

A chemist produced 11.9 g of aluminum, Al.  How many moles of aluminum were produced?

Step 1:  Given ___11.9 g Al_____________      Unknown mol Al_______

Step 2:   

How many moles of silver, Ag, are in 3.01 x 10²³ atoms of silver?

Step 1:  Given ___3.01 x 10²³ atoms Ag___________      Unknown mol Ag

Step 2:   

What is the mass in grams of 1.20 x 10⁸ atoms of the element copper, Cu?

Step 1:  Given ___1.20 x 10⁸_____________      Unknown ____g Cu_______

Step 2:   

Step 3: 1.27 x 10^-14

• How many atoms of carbon are in 0.020 g of carbon?

1.204 x 10²²/12.0

• What is the mass in grams of 7.5 x 10¹⁵ atoms of nickel?

Nuclear Reactions

Nuclear Reactions

• A reaction that changes the nucleus of an atom

  • Ways to change

    • Gain or lose protons

    • Gain or lose neutrons

• Nucleons: A general term that encompasses protons and neutrons

• Transmutation: A change in the identity of a nucleus as a result of a change in the number of protons or neutrons

Transmutation

• Artificial Transmutation: Bombardment of nuclei with charged particles

  • Not naturally found in nature

  • Neutrons used because there is no repulsion from other particles

• Transuranium Elements: Elements with more than 92 protons in their nuclei

  • All the elements passed Uranium don’t have a stable nucleus and use Uranium to be created, it does not exist in nature

  • None are naturally occurring (found in nature)

  • They decay very fast and are made by humans

Writing Elements in Nuclear Chemistry

• Mass number as the superscript

• Atomic number as the subscript

• Symbol of the element

(Protons on the bottom, mass of the element on the top)

• Radioactive decay: is the spontaneous disintegration of a nucleus into a slightly lighter nucleus

• This happens when emission of particles and/or electromagnetic radiation

Particle Types

• Quarks effect this because the nucleus can loose the equivalent of an electron from the nucleus, a neutron turns into a proton, 1 more neutron, change what element you have (1 atomic number bigger than it previously was

• The positive version of a beta particle, positive charge, but no mass, turns a proton into a neutron, loose a positive charge so charge is 1 lower

• Energy, has no mass or charge of protons in the nucleus, does not affect what the nucleus looks like, release a lot of dangerous energy that can have a lot of bad short/long term effects on humans

Radiation Types

• Alpha Charge

• Beta Charge

• Positron

• Gamma Radiation

Writing Radioactive Decay Equations

• The total mass number and atomic number must be equal on both sides of the equation

• Determine the total mass number and atomic number for each side of the reaction

• Add/subtract the mass number or atomic number to make the sides equal.  Be sure to show what types of particles are given off during the decay

208  Pb (Lead)

82

• In the late 1700s Antoine Lavoisier compiled a list of all the unknown elements.

• Lavoisier was killed during the French Revolution in 1794

• John Newlands noticed that when the elements were arranged by increasing atomic mass, their properties repeated every eighth element (The Law of Octaves)

Dimitri Mendeleev and Lothar Meyer 

• Mendeleev predicted the existence and properties of unknown elements, leaving blank spaces in the table where he thought the undiscovered elements should go.

Henry Moseley

• Arranged elements by increasing the number of protons, rather than mass.

• Periodic Law: There is a periodic repetition of chemical and physical properties of elements when they are arranged by increasing atomic number

The Modern Periodic Table

• Columns in the periodic table are referred to as groups.

• Rows in the periodic table are called periods. 

Metals vs. Nonmetals

Metals tend to be:

• Lustrous: shiny

• Malleable: able to be hammered into a thin sheet

• Ductile: able to be drawn into a wire

• Good conductors of heat and electricity

• Able to lose electrons to attain a complete valence shell

• Unreactive with each other (instead they mix to form alloys)

Non-Metals tend to be:

• Dull

• Brittle

• Poor conductors of heat and electricity

• Able to gain electrons to attain a complete valence shell

• Reactive with each other to form molecular compounds

Metalloids

• The elements bordering the metals and nonmetals are called metalloids

• Metalloids have properties of both metals and nonmetals (somewhat shiny, malleable, ductile, conductive…)

Alkali Metals

• The alkali metals are all shiny, soft, highly reactive metals at standard temperature and pressure 

• They have 1 electron in their outermost valence shells.

• They readily lose their outermost electron to form cations with charge +1.

Alkaline earth metals

• The alkaline earth metals are all shiny, silvery-white, somewhat reactive metals at standard temperature and pressure.

• They have 2 electrons in their outermost valence shells.

• These elements readily lose those 2 electrons to form cations with charge +2

Transition metals

• The transition metals form a large range of complex ions with various charges

• They tend to be highly coloured complexes with catalytic properties either as the element or as ions (or both)

• They are characterized by having d electrons in their valence shell.

Halogens

• Halogens have seven valence electrons in their outermost energy level, so they will easily gain an electron by reacting with atoms of other elements to attain a filled valence shell of electrons.

• Halogens exist in nature as salts or diatomic molecules

Noble gasses

• Noble gases are odorless, colorless, monoatomic gases that have a filled outer shell of electrons.

• They are very stable and generally do not react with any other elements.

Lanthanides and Actinides

• The lanthanides and actinides are all soft metals and many are radioactive. 

• Most of the lanthanides can be found naturally on Earth, but the actinides are typically made in nuclear reactors and not found in nature.

Atomic Radius

• The atomic radius of a chemical element is a measure of the size of its atoms, usually the average distance from the center of the nucleus to the boundary of the surrounding cloud of electrons.

Atomic Radius

• What is the trend going across a period? 

• Going left to right, atoms get smaller

• What is the trend going down a group?

• Down the periodic table, atoms get bigger

Ionization energy: The ionization energy is the energy required to remove an electron from a gaseous atom (the energy required to form a cation)

• How strongly does the atom’s nucleus hold on to its valence electrons?

• What is the trend going across a period?

• As you go across a period, it gets larger

• What is a trend going down a group?

• Ionization energy decreases, the further down the periodic table you go, the harder it is to give away an electron

• The smaller the size, the larger the ionization energy

• Big atom, it's easier to take electrons away

Electron Affinity

• The electron change that occurs when an electron is added to a gaseous atom is called an electron affinity.

• For most atoms, energy is released when an electron is added. 

• The greater the attractions between the atom and the added electron, the more energy is released.

Cl(g) + e⁻ → Cl⁻(g) ΔE = − 349 kJ/mol

Ar(g) + e⁻ → Ar⁻(g) ΔE = > 0 kJ/mol

• Big negative number, really wants electrons

• Chlorine has a high electron affinity bc of big negative number

• Argon has a low electron affinity bc of low negative number or positive number

• What is the trend going across a period?

• Electron affinity increases across a period, going left to right they are more likely to want an electron

• What is a trend going down a group?

• Tends to be pretty similar within a group

Electronegativity

• The electronegativity of an element indicates how strongly the element attracts electrons towards itself in a chemical bond. 

• How much does the atom want to gain another electron?

• Electronegativity is measured on the Pauling scale, 0-4.

• What is the trend going across a period?

• Electronegativity increases, towards the right, they are more likely to steal electrons

• What is a trend going down a group?

• Going down a group, electronegativity gets smaller

• Top right, has highest electronegativity (F)

• Ionization and electronegativity go together

• Ex. Small atom has high electronegativity and ionization energy

Density

• What is density? What is the formula for density? M/V

• Can you predict the density of germanium based on the periodic trend going down a group?