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MSUB / Chemistry / CHMY 121 / Give examples of physical properties/changes.

Give examples of physical properties/changes.

Give examples of physical properties/changes.

Description

School: Montana State University - Billings
Department: Chemistry
Course: Intro to General Chemistry
Professor: Rhonda dillman
Term: Spring 2017
Tags: General Chemistry, Chemistry, ChemistryLecture, Study Guide, exam, exam1, conversion, Studyquestions, Chem, and General Chemistry notes hand written examples walk through easy difficult Scientific Notation/Components of Numbers Counting Significant Figures and Calcula
Cost: 50
Name: Chemistry Exam 1 Study Guide
Description: This study guide includes: - Exam information - Terms and Definitions List - Study quizzes - Formulas - Metric prefixes - Conversion sheet. - Detailed and color-coded lecture notes
Uploaded: 01/30/2017
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CHMY 121-001 Exam 1 Study Guide


Give examples of physical properties/changes.



Chapter 1:

Key Terms and Definitions, Numbers, Units,Formulas

(1-20-17)  

Chemistry: The study of matter.  

Matter: Anything that has mass and occupies space.  

Light: A form of energy (not matter), in addition to heat

In this course, we are going to discuss:

- The structure of matter

- The properties of matter  

- The reactions of matter

Scientific Method: 

- We make an observation

- We make a hypothesis  

- We make experiments to either back up or negate the hypothesis

- We state a theory (which can be disproven)

Solids: 

- Defined as having a definite shape and a definite volume 

Liquids: 

- Defined as having a definite volume but an indefinite shape 

- They take on the shape of any container you put them into

Gases: 

- Defined as having an indefinite volume and an indefinite shape 

- They take on both the volume and the shape of the container you put them in.

Physical properties/changes in matter:

- Properties that we can observe without changing the chemical makeup of matter.  

- Changes in the physical states of the matter are known as physical properties.

- Can be used to help identify a substance

o Examples:

▪ Matter changing between a solid state, a liquid state, and a gaseous state through melting  points and boiling points

▪ Color

▪ Density


How will you define pure substance?



Don't forget about the age old question of Describe carbohydrates?
Don't forget about the age old question of What are homologous characteristics?
We also discuss several other topics like What is the meaning of rock in geology?

Chemical properties/changes in matter:

- A chemical change involves a change in the chemical makeup of a substance  

- A substance will never return to its original form after a chemical change has taken place - Can be used to help identify a substance If you want to learn more check out Who is mary calkins?

o Examples:

▪ Rust forming on Iron after the Iron reacts with Oxygen  

▪ Burning anything, because it combusts molecules (burning toast, for instance)

▪ Leaves changing color  

The two groups of matter: 

Pure Substances: 

- Pure Substances have constant compositions and constant properties 


Enumerate the 3 characteristics of subatomic particles.



Two groups of pure substances: 

- Elements: 

o The simplest type of pure substance in existence – building blocks of all matter

o Elements cannot be broken down into anything simpler, even by chemical means  

o Constant compositions – composed of identical atoms

o Elements can be found on the periodic table, each with a symbol associated with them.  ▪ One or two-letter symbols – the first letter is always capitalized and the second letter is always  lowercase  

o The colors of the symbols are significant Don't forget about the age old question of What is the meaning of genetics in biology?

▪ Black lettering = the element is solid at room temperature  

▪ Blue lettering = the element is liquid at room temperature

▪ Red lettering = the element is a gas at room temperature

▪ = the element is radioactive 

o Diatomic elements:

▪ Always occur as a result of two atoms bonding to each other – never individual atoms  ▪ There are 7 diatomic elements: Bromine, Iodine, Nitrogen, Chlorine, Hydrogen, Oxygen, and  Fluorine.  

∙ Strategy to remember the elements: Have No Fear Of Ice Cold Beer.

- Compounds:  

o A chemical combination of two or more elements 

o Constant, non-varying compositions and properties If you want to learn more check out What is the function of neurons?

o A chemical formula tells us what elements are present in a compound, and how many atoms of each  element is present. The atoms present are represented by a subscript.  

Mixtures: 

- The physical combination of two or more pure substances

- Mixtures vary in compositions and properties 

- Mixtures don’t change the chemical makeup of substances, you can separate substances back out again through  physical means into their individual components

- The composition of mixtures can vary

o With varying composition comes varying properties.

Two groups of mixtures: 

- Homogenous mixtures: 

o Only one phase present – you can’t tell by looking at it that there are two or more pure substances present - Heterogeneous mixtures: 

o You can physically see two or more phases (or pure substances) present

Qualitative and Quantitative Measurements: 

Qualitative measurements: 

- Subjective  

- No number or unit is associated with them

- Examples:

o “Big dog”

o “Small flower”  

- Cannot be used effectively in chemistry  

Quantitative measurements: 

- There is a number and a unit associated with them. In this class, every number will have a unit associated with it. - Units describe what kind of measurement you’re taking

Metric system prefixes: 

Kilo (k) 

- 1000  

Centi (c) 

- 1/100th 

Milli (m) 

- 1/1000th 

Micro (μ) 

- 1/1000000th

(1-23-17)  

Length measurements: 

- Metric system measurement for length: meter (m) 

o Kilometer (km): 1000 meters 

o Centimeter (cm): 1/00th of a meter 

o Millimeter (mm): 1/1000th of a meter 

o Micrometer (μm): 1/1000000th of a meter 

Mass/weight measurements: 

- Metric system measurement for mass: gram (g) 

o Kilogram (kg): 1000 grams 

o Milligram (mg): 1/1000th of a gram 

o Microgram (μg): 1/000000th of a gram 

Volume measurements: 

- Metric system measurement for mass: liter (L) 

o Kiloliter (kL): 1000 liters 

o Milliliter (mL): 1/1000th of a liter 

o Microleter (μL): 1/1000000th of a liter 

Counted number vs. measurement: 

Counted numbers: 

- Have no uncertainty to them  

Measurements: 

- Always have some uncertainty associated with them

- Two reasons for uncertainty:

o The instrument used to make the measurement

o The person making the measurement  

We can’t read more accuracy into a measurement than the instrument allows for.  

- Read exactly what you can off of an instrument, then estimate past that  - The last place is the degree of uncertainty 

- Examples:

o 1.06 = The degree of uncertainty is +or- .01 

o 15 = The degree of uncertainty is +or- 1 

o 1.1 = The degree of uncertainty is +or- .1

Significant Figures (also known as ‘sigfigs’ or ‘SF’)

- Within a measurement, all non-zero digits are significant  

- Leading zeros are never significant (leading zeros come before any non-zero digit)  - Confined zeros are always significant (confined zeros are between non-zero digits) - Trailing zeros may or may not be significant

1. Trailing zeros are significant if there is a decimal point in the number, and if there is  no decimal point, trailing zeros are not significant

2. A bar over the top of a trailing zero means that it’s significant  

Significant Figures in calculations:  

- Multiplication and division:  

1. The answer must have the same amount of significant figures as the measurement  with the least number of significant figures  

- Addition and subtraction:

1. The answer must go to the place that is common to all of the measurements  

- Never use counted numbers to determine the amount of significant figures in your answer - Never use defined numbers to determine the amount of significant figures in your answer 1. Defined numbers are numbers you looked up in a book, not a measurement that you  yourself made

- Only use the measurement that you made to determine the amount of significant figures Scientific Notation: 

- Makes really big or small numbers easier to write  

- If the decimal moves to the right, the exponent is negative. If the decimal moves to the left,  the exponent is negative

- If there is a negative exponent, that means that the original number is less than one - If there is a positive exponent, that means that the original number is greater than one  

Dimensional Analysis: 

- Also called the ‘unit conversion method’ 

- Relies on conversions and ratios  

- When working through the problem, you must ask yourself:

1. What do you want?

2. What do you need?

- When working through dimensional analysis problems in exams, always show all of your  work and the steps you took to get your answer, otherwise the answer could be marked  incorrect.

(1-27-17)

Density: 

- A physical property of a substance (you can observe it without changing the chemical composition of  the compound)

- The mass of a substance per unit volume  

o Density = mass/unit volume

o d=m/V 

- Units of density:  

o The units vary depending on if the substance is solid, liquid, or gas

o Solid:

▪ Gram/cubic cm  

▪ g/cm^3 

o Liquid:

▪ Gram/milliliter

▪ g/mL 

o Gas:

▪ Gram/liter

▪ g/L 

- You can distinguish between substances by measuring their densities  

- For example:

o Water typically has the density of 1 gram per mL(1 g/mL)

Never worry about significant figures and rounding until you get your final answer  

(1-30-17)

Specific Heat: 

- Specific heat is a physical property  

- Specific heat deals with energy and temperature 

Temperature: 

- Temperature is an indicator of heat 

- Temperature is measured with three different scales:

o Fahrenheit (which is ignored in this class)

o Celsius  

o Kelvin 

- The only difference between Celsius and Kelvin is the freezing point/boiling point of water - 0 K is absolute zero, which is where all molecular motion stops

- Freezing point of water:

o 0 degrees C 

o 273 K 

- Boiling point of water:

o 100 degrees C 

o 373 K 

- Temperature and significant figures:

o In temperature, trailing zeros without a decimal point are considered significant  ▪ This is true for temperature only 

- Because temperature is an indicator of heat, it’s really an indicator of energy, because heat is a form  of energy  

Energy: 

- The capacity to do work or to supply heat 

- One form of energy is heat, another form of energy is light

o The indicator of this heat is temperature  

- Energy can be measured in units of calories (cal) 

o A calorie is the amount of heat it takes to raise the temperature of one gram of water by one  degree C

Specific Heat (SH) 

- The amount of heat it takes to raise the temperature of 1 gram of any substance by 1 degree C - Units:

o cal/(g)(C) 

- Specific Heat Formula:

o The amount of energy divided by the mass of the substance in grams times the change of  temperature in degree C

o SH = cal/mass* 

T

▪ The change in temperature is always positive  

- Specific Heat can be used to identify a compound  

- Because of the units, in any answer, the mass has to be in grams, and the temperature has to be in  Celsius

- ‘Cal’ (with a capital C) is a nutritional calorie, and is equal to 1000 calories (cal, with a lowercase c) o Cal = kcal

(2-1-17)

Atom: 

- The smallest particle of an element that still retains the properties of the element 

- An atom has to be whole in order to retain the properties of the element  

- An atom is always neutral  

Subatomic Particles: 

- There are 3 kinds of Subatomic Particles:

o 1. Proton (p) 

o 2. Neutron (n) 

o 3. Electron (-) 

3 characteristics of subatomic particles:

- Mass:

o Protons and neutrons have the same mass

▪ 1 atomic mass unit (amu)  

o Electrons have a mass that is 10^4 times smaller than a proton or a neutron  

▪ Compared to a proton and neutron, we say that electrons have no mass (0 amu)

- Charge:

o Protons and electrons have equal but opposite charges  

▪ Protons: +1

▪ Electrons: -1  

o Neutrons are neutral and have no charge  

- Location within the atom:

o Protons and neutrons are both found in the nucleus (the center of the atom) 

o Electrons are found in the space outside of/around the nucleus, but they are never in the nucleus Nucleus: 

- Has a positive charge, because of the protons inside it  

o The space around the nucleus has a negative charge

- The nucleus is very dense compared to the space around it  

- All of the mass of an atom is in the nucleus  

Because an atom is neutral, the number of protons and electrons have to be equal to each other  Atomic Number: 

- The atomic number is the number on the top of an element in the periodic table

- The atomic number gives us the number of protons in the atom of an element  

o The number of protons can never vary

▪ You can always identify an element by the number of protons in the atom  

- The atomic number can be used to identify elements

- Because the amount of protons and electrons in an atom must be equal, we can identify the amount of electrons,  as well

o However, the number of electrons can vary, so the number of electrons can’t be used to identify an  element  

o The number of electrons in an atom tells us the chemical behavior of an element

Mass Number: 

- The number below the symbol of the element on the periodic table is the mass number

- The mass number gives us the number of protons plus the number of neutrons  

- The mass number is the weighted average of all the isotopes of a given element  

o Weighted average:

▪ Equal to the sum of the 1st isotope percentage times its mass plus the 2nd isotope percentage times  its mass 

- If you subtract the atomic number from the mass number, you get the number of neutrons in the element  o Mass number – Atomic number = number of neutrons 

The number of neutrons in the atom of an element can vary

- Atoms of a given element can vary in their number of neutrons  

- The variation in neutrons gives us the isotopes of an element

Isotopes: 

- Atoms of a given element that vary in their number of neutrons

The Periodic Table:

- Elements are arranged by increasing atomic number, and elements with similar chemical behaviors with occur at  periodic intervals  

o Chemical behaviors occur due to the similarities between the electrons of the atoms of the elements - Elements in the vertical columns have similar chemical behavior

o The vertical columns are called groups 

o Electrons determine the chemical behavior

- Elements in the horizontal rows are called periods 

Groups: 

- The tall groups are the ‘A’ groups 

o All of the A group elements are referred to as the ‘representative elements’ 

▪ They are called the ‘representative elements’ because they behave the way they’re supposed to – they ‘represent’ the rules

- The short groups are the ‘B’ groups 

o They are the transition metals 

- Everything to the left of the ‘red stair step line’ is a metal, and everything to the right is a non-metal Metal vs. Non-metal:

- Metal: 

o All solid at room temperature (except Mercury)

o Shiny

o Shapeable metals

o Conduct electricity and heat  

o Metal atoms like to lose electrons in chemical reactions  

o High melting points (melting point is above 300 C)

- Non-metal  

o Range from being solid, liquid, or gas at room temperature  

o Dull, brittle, non-shapeable  

o Doesn’t conduct heat and electricity  

o Atoms of non-metals like to gain electrons in chemical reactions

o Low melting points (melting point is less than 300 C)

Halogens: 

- 7A group of elements

Inert/Noble Gases: 

- 8A group of elements  

Group and period designations can be used to determine what an element is on the periodic table  

- Example:

o Sulfur (S): Period 3, Group 6A

An element’s placement on the periodic table tells us how the element reacts chemically

(2-3-17)

Electron arrangements in an atom:

An electron’s location in an atom is based on the energy that the electron has.  

- Essentially, there are different energy levels within an atom  

Shell: 

- The area around the nucleus that contains electrons of a given energy  

- Each shell in an atom contains a different level of energy  

- Shells are designated with numbers

o 1, 2, 3 and so on  

o Numbered from the shell closest to the nucleus, moving outward  

- Shell 1 (closest to the nucleus) has the lowest amount of energy and the energy increases from there  - The maximum number of electrons per shell is equal to 2n^2, where n is equal to the shell number  o Max #e/shell = 2n^2  

o N = shell#

▪ Shell 1 max electrons = 2 

▪ Shell 2 max electrons = 8 

▪ Shell 3 max electrons = 18 

Subshells: 

- An area within a shell that contains electrons of a given energy  

- Designated with letters:

o The subshells are s, p, d, and f 

o Energy increases from s to f within a shell  

▪ s has the lowest amount of energy, f has the highest

- Not every shell contains every subshell

- The number of subshells within a shell is equal to the shell number  

o # of subshells/shell = the shell #  

▪ Shell 1 = s  

▪ Shell 2 = s and p  

▪ Shell 3 = s, p, and d  

▪ Shell 4 = s, p, d, and f

- The maximum number of electrons within a subshell is dependent on how many orbitals a  subshell has  

Orbitals: 

- Areas in a subshell that contain electrons of a given energy 

- The maximum number of electrons in any orbital is 2  

o Max # e/orbital = 2 

- Within a subshell, the orbitals all have equal energy  

- Subshell s = 1 orbital, 2 electrons 

- Subshell p = 3 orbitals, 6 electrons 

- Subshell d = 5 orbitals, 10 electrons 

- Subshell f = 7 orbitals, 14 electrons 

Electrons always go to the lowest energy level first  

Electron configuration of an atom: 

- Shows where the electrons are within an atom 

- The number of electrons in a subshell is written as a superscript  

- Full electron configurations of the first 11 elements:

o Hydrogen = 1s^1

o Helium = 1s^2  

o Lithium = 1s^2 2s^1  

o Beryllium = 1s^2 2s^2  

o Boron = 1s^2 2s^2 2p^2  

o Carbon = 1s^2 2s^2 2p^2

o Nitrogen = 1s^2 2s^2 2p^3  

o Oxygen = 1s^2 2s^2 2p^4  

o Fluorine = 1s^2 2s^2 2p^5  

o Neon = 1s^2 2s^2 2p^5  

o Sodium = 1s^2 2s^2 2p^6 3s^1  

- The period number that an element is in tells us the highest numbered shell in that atom that  contains electrons  

- s contains a maximum of 2 electrons  

- s groups

o Elements in group 1A end in an s^1 configuration

o Elements in group 2A end in an s^2 configuration  

- p groups

o All other A groups end in a p configuration  

- 3d is higher in energy than 4s, so 4s fills with electrons before 3d does  

o d is always one behind

- Abbreviated electron configuration:  

o Abbreviated electron configuration is based on the noble gas that precedes the element o Examples of abbreviated electron configuration:

▪ Silver (Ag) = [Kr] 5s^2 4d^9  

▪ Phosphorus (P) = [Ne] 3s^2 3p^3

- Orbital diagrams: 

o Orbital diagrams show the orbitals we put electrons into  

o The lowest energy is at the bottom and it increases in energy moving upwards

o One electron has to go in each orbital before they can begin to be paired  

o Example of an orbital diagram:

Valence Electrons: 

o The electrons in the highest numbered shell are the valence electrons  

o The group numbers for the A Groups tells you how many electrons are in the highest numbered  shell  

o A Group elements:

▪ Group # = # e in the highest # shell

Formulas: 

Density:  

- d=m/V

- V=m/d

- m=d*V

Specific Heat: 

- SH = cal/(mass)(ChangeinTemperature)

- cal = (SH)(mass)(ChangeinTemperature)

- mass = cal/(ChangeinTemperature)(SH)

- Final Temperature = (cal/(mass)(SH)) + initial temperature  

Mass number: 

- Atomic Weight = ∑[(isotope% x mass)]

Electrons per shell: 

- Max #e/shell = 2n^2

For the exam:

Memorize one conversion factor to get you from the English system to the metric system Memorize what the conversions among the metric system are  

During the exam, you can ask the professor for English to English conversions, but she won’t give you  metric system conversions

Memorize the formulas needed (the initial formula if nothing else, because that formula can be converted  from there)

Know all three kinds of electron configuration  

Study Methods:

- Work through problems at the end of chapters 1 and 2 in the book  

- Do the ‘dynamic study modules’ on the website where you find the online homework  o The “Pearson Dynamic Study Modules” app can be downloaded on smartphones for free,  this lets you work through short study quizzes on your phone.  

- I’ve created a Quizlet folder for this course that can be studied from:

o https://quizlet.com/class/4000037/

Study Quizzes 

Identify if the term is a pure substance or mixture 

1. Sodium  

a. Pure substance

b. Mixture

2. Water  

a. Pure substance

b. Mixture

3. Soil  

a. Pure substance

b. Mixture

4. Coffee

a. Pure substance

b. Mixture

5. 70% isopropyl alcohol  

a. Pure substance

b. Mixture

6. Salt Water  

a. Pure substance

b. Mixture

7. Carbon Dioxide

a. Pure substance

b. Mixture

8. Table salt  

a. Pure substance

b. Mixture

9. Air

a. Pure substance

b. Mixture

10. Milk

a. Pure substance

b. Mixture

11. Atmospheric Gas

a. Pure substance

b. Mixture

12. Sugar

a. Pure substance

b. Mixture

13. Baking Soda  

a. Pure substance

b. Mixture

Answers:

1:A. 2:A. 3:B. 4:B. 5:B. 6:B. 7:A. 8:A. 9:B. 10:B. 11:B. 12:A. 13:A

Identify if the pure substance is an element or a compound  

1. Sodium  

a. Element  

b. Compound

2. Water

a. Element  

b. Compound

3. Oxygen  

a. Element  

b. Compound

4. Carbon Dioxide  

a. Element  

b. Compound

5. Table salt  

a. Element  

b. Compound

6. Sugar  

a. Element  

b. Compound

7. Baking soda  

a. Element  

b. Compound

8. Gold  

a. Element  

b. Compound

9. Sulfuric Acid

a. Element  

b. Compound

10. Iron

a. Element  

b. Compound

Answers:

1:A. 2: B. 3:A. 4:B. 5:B. 6:B. 7:B. 8:A. 9:B. 10:A.

Identify if the mixture is homogeneous or heterogeneous  

1. Soil  

a. Homogeneous  

b. Heterogeneous

2. Coffee

a. Homogeneous  

b. Heterogeneous

3. 70% Isopropyl Alcohol  

a. Homogeneous  

b. Heterogeneous

4. Salt Water

a. Homogeneous  

b. Heterogeneous

5. Cake Batter  

a. Homogeneous  

b. Heterogeneous

6. Air  

a. Homogeneous  

b. Heterogeneous

7. Pond Water

a. Homogeneous  

b. Heterogeneous

8. Meat

a. Homogeneous  

b. Heterogeneous

9. Juice  

a. Homogeneous  

b. Heterogeneous

10. Milkshake  

a. Homogeneous  

b. Heterogeneous

11. Mud

a. Homogeneous  

b. Heterogeneous

12. Sand  

a. Homogeneous  

b. Heterogeneous

13. Water with Oil

a. Homogeneous  

b. Heterogeneous

14. BBQ Sauce

a. Homogeneous  

b. Heterogeneous

15. Soda  

a. Homogeneous  

b. Heterogeneous

Answers:

1:B. 2:A. 3:A. 4:A. 5:A. 6:A. 7:B. 8:B. 9:A. 10:B. 11:B. 12:B. 13:B. 14:A. 15:A

Identify how many significant figures a number has 

1. 203.0  

a. 1 SF  

b. 2 SF

c. 3 SF

d. 4 SF  

2. 2030  

a. 1 SF  

b. 2 SF

c. 3 SF

d. 4 SF  

3. 29.35

a. 1 SF  

b. 2 SF

c. 3 SF

d. 4 SF

4. 1.230

a. 1 SF  

b. 2 SF

c. 3 SF

d. 4 SF

5. 1230

a. 1 SF  

b. 2 SF

c. 3 SF

d. 4 SF

6. .0950

a. 1 SF  

b. 2 SF

c. 3 SF

d. 4 SF

7. .006

a. 1 SF  

b. 2 SF

c. 3 SF

d. 4 SF

Answers:

1:D. 2:C. 3:D. 4:C. 5:C. 6:C. 7:A.

Conversion Sheet:

Length: 

- 1 km = 1000 m

- 1 m = 100 cm  

- 1 cm = 10 mm

- 1 m = 1000 mm

- 1 km = 0.6214 mi

- 1 km = 3280.84 ft

- 1 in = 2.54 cm  

- 1 ft. = 30.48 cm  

- 1 yd = 91.44 cm

- 1 mi = 1.6 km

- 1 ft = .3048 m  

- 1 ft = .0003048 mm

Voume: 

- 1 L = 1000 mL  

- 1 L = .26 gal.  

- 3.79 L = 1 gal.  

- .96 L = 1 qt.  

- 1 L = 1.06 qts.

- 1 mL = 0.00211338 liquid pints

- 1 mL = 1 cm^3 = 1cc

- 29.6 mL = 1 fl. oz.  

- 1 L = 1000000 uL

- 1 cup = 0.236588 L

- 1 cup = 236.588 mL

Mass/Weight: 

- 1 kg = 1000 g

- 1 g = 1000 mg  

- 1 lb. = 454 g  

- 1 oz. = 28.4 g  

- 1 kg = 2.20 lb

- 1000000 ug = 1 gram

Converting Temperature:

Celsius to Fahrenheit and Kelvin:

F = (1.8)( C ) + 32.0

K = C + 273.15

Fahrenheit to Celsius and Kelvin:

C = (F – 32.0) / 1.8

K = (5/9)(F + 459.67)

Kelvin to Celsius and Fahrenheit:

C = K – 273.15

F = (9.5)(K) – 459.67

Exam information: 

- Multiple choice  

- Includes definitions of terms

- She’ll give the symbol of the elements and you have to spell the correct name of the element,  and vice versa

- There is a section where you have to do a full electron configuration, abbreviated electronic  configuration, and orbital diagrams

- Calculations portion - density, specific heat, and dosage problems. Expect conversions in all of  the problems

○ Know at least one English to metric conversion  

○ Know all metric prefixes  

○ Bring calculator!  

- Never leave anything blank  

○ At least write the equations, you get points for knowing that  

- Show all of your work  

- The answer will be marked wrong if you do not show your work

- Know what valence electrons are  

- Know the transition metals, non-metals, metals, representative elements, etc.

Terms and Definitions: 

Chemistry:

- The study of matter

Matter:  

- Anything that has mass and occupies space  

Light:

- A form of energy  

Solids, Liquids, and Gases:

- Solid: Has a definite shape and a definite volume  

- Liquid: Has a definite volume but an indefinite shape  

- Gas: Has an indefinite volume and an indefinite shape  

Physical properties/changes:

- Properties that we can observe without changing the chemical makeup of matter

Chemical properties/changes:

- A chemical change involves a change in the chemical makeup of a substance

Pure Substance:

- Have constant compositions and constant properties  

Elements:

- The simplest type of pure substance in existence – the building blocks of all matter. Cannot be broken  down into anything simpler  

Diatomic elements:

- Always occur as a result of two atoms bonding to each other – never individual atoms Compounds:

- A chemical combination of two or more elements. Constant compositions and properties Chemical Formula:

- Shows what elements are present in a compound, and how many atoms of each element is present

Mixture:

- The physical combination of two or more pure substances, varying compositions and properties Homogenous mixture:

- Only one phase present – you can’t tell by looking at it that there are two or more pure substances  present  

Heterogeneous mixture:

- You an physically see two r more pure substances present

Qualitative measurements:

- Subjective; no number or unit is associated with them  

Quantitative measurements:

- There is a number and a unit associated with them; units describe what kind of measurement you’re  taking  

Density:  

- A physical property, density equals the mass of a substance per unit volume  

Temperature:

- An indicator of heat

Energy:

- The capacity to do work or to supply heat. One form of energy is heat, another form of energy is light  Calorie:

- The amount of heat it takes to raise the temperature of one gram of water by one degree Celsius Specific Heat:  

- The amount of heat it takes to raise the temperature of 1 gram of any substance by 1 degree Celsius. A  physical property; deals with energy and temperature.  

Atom:

- The smallest particle of an element that still retains the properties of the element. Always neutral.  

Nucleus:

- The center of the atom, has a positive charge, and very dense compared to the space around it. All of  the mass of an atom is in the nucleus

Atomic Number:

- The number on the top of an element in the periodic table – represents the number of protons in the  atoms of an element  

Mass Number:

- The number below the symbol of the element on the periodic table – represents the number of protons  plus the number of neutrons.

Isotopes:

- Atoms of a given element that vary in their number of neutroms  

Groups and Periods:

- Groups: The elements in the vertical columns of the periodic table  

- Periods: The elements in the horizontal rows of the periodic table  

Representative Elements:

- The A Groups

Transition Metals:

- The B Groups  

Metal:

- Except Hg, solid at room temperature. High melting points. Always lose electrons in chemical  reactions  

Non-metal:

- Can be solid, liquid, or gas at room temperature. Low melting points. Always gain electrons in  chemical reactions.  

Halogens:

- 7A group of elements  

Inert/Noble Gases:

- 8A group of elements  

Shell:

- The area around the nucleus that contains electrons of a given energy  

Subshells:  

- An area within a shell that contains electrons of a given energy

Orbitals:

- Areas in a subshell that contain electrons of a given energy – the maximum number of electrons in any  orbital is 2  

Electron configuration of an atom:

- Shows where the electrons are within an atom  

Valence Electrons:

- The electrons in the highest numbered shell; the group numbers for the A Groups represent how many  electrons are in the highest numbered shell.  

Mass:  

- A measure of the amount of matter in an object

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