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AU / Chemistry / CHEM 1030 / What are qualitative properties?

What are qualitative properties?

What are qualitative properties?

Description

School: Auburn University
Department: Chemistry
Course: Fundamentals Chemistry I
Professor: John gorden
Term: Fall 2015
Tags: CHEM 1030 and BRETT CAGG
Cost: 50
Name: CHEM 1030 Cagg Exam 1 Study Guide
Description: Includes a list of vocab words (ch 1-3.4), practice problems, important formulas, a set of notes from both chapters 1 and 2!
Uploaded: 02/06/2016
21 Pages 96 Views 7 Unlocks
Reviews


CHAPTERS 1-3.4 VOCABULARY


What are qualitative properties?



• Solid: particles tightly packed with no space in between (ice) • Liquid: particles are packed loosely with some space in between  (water at room temperature)

• Gas: particles are not packed and roam freely (water vapor) • Homogenous mixtures: these mixtures are thoroughly mixed to  where you cannot tell the original substances (also called a solution). • Heterogeneous mixtures: these are not thoroughly mixed, and you  can often tell what the original substances were.

• Quantitative properties: are measured or expressed using a number,  ALWAYS ADD UNITS!  

• Qualitative properties: are based mainly on observation  • Physical properties: are those, which can be observed/measured  without changing the identity of the substance. Some types of  physical properties are: color, melting point, boiling point, smell, etc. • Physical changes: are those, which disrupt the state of matter of a  substance (the substance PHYSICALLY changes).


What are physical properties?



If you want to learn more check out What is the function of serous membranes?
We also discuss several other topics like Where do the other amino acids exist?

• Chemical properties: is a reaction that a substance has when  interacting with another substance.

• Chemical changes: reactions that result in the total change of the  composition of the substance, so that the original one does not exist  anymore.

• Extensive properties: depend on how much matter there is in an  object, the more matter an object has the more its mass is  • Intensive properties: don’t depend on how much matter there is in  an object, temperature and density are 2 types

• Mass: the measure of the amount of matter in an object  • Weight: the force which is exerted by an object due to gravity, the  weight of an object varies from place to place since gravity is  involved


What are chemical properties?



We also discuss several other topics like Who can also set price floors (minimum prices) and price ceilings (maximum prices) on certain goods?

• Significant figures: are the MEANINGFUL digits in a number • Accuracy: shows how close or exact a measurement is to the real  value  

• Precision: shows how close a set of measurements are to one  another  

• Conversion factors: fractions where 2 quantities that mean the same  thing (equal to each other) are put one in the numerator and other  in the denominator. We also discuss several other topics like What is cognitive psych?

• Dimensional analysis: also known as the factor label method, is a way  of converting units in problem solving  

• Atoms: are the building blocks and smallest units of matter • Element: contain a unique form of the same atom (each element has  a different form of atoms), which cannot be divided into 2 or more  simpler substances

• Radiation: is the transmission of energy in the form of waves  • Cathode ray tube: also known as the electron beam, is a glass tube  with metal plates on each end, which has all air, sucked out from it.  This device detects electrons using radiation.

• Cathode: is a negatively charged plate within the cathode ray which  produces radioactive waves called cathode rays

• Anode: is a positively charged plate within the cathode ray • Columb’s law: like charges push against one another and opposite  charges come towards one another Don't forget about the age old question of What is the relationship between the numerator in the traditional column and the numerators in the refined column?

• X-rays: rays that are able to penetrate matter and cause materials to  give off a fluoresce light. They do no have charged particles due to  them not being deflected by anything.

• Radioactivity: a spontaneous emission of rays that are highly  energetic and cannot be deflected due to anythingIf you want to learn more check out What type of landscape conveys a sense of power?

• Alpha α rays: have positively charged particles, which are called  alpha particles. These are deflected away from a positively charged  plate.  

• Beta β rays: have negatively charged particles called beta particles  (also known as electrons). These are deflected away from negatively  charged plates.

• Gamma y rays: these have no charge

• The plum pudding model: an idea proposed by Thomson saying that  electrons were embedded into the atom like “raisins in a scoop of  rum ice cream”; it was the accepted theory for quite a while.

• The atomic number: the number that defines the element, it also  represents the number of protons that are in the nucleus of that  atom.

• The atomic mass number: the total number of protons and neutrons  that are in the nucleus of an element’s atom.

• Isotopes: are atoms, which have the same atomic number, but  different atomic mass, (the number of neutrons is different from the  number of protons).

• Atomic mass: the mass of an atom in atomic mass units (amu) • The periodic table: consists of 118 elements that are grouped  according to their physical and chemical properties.

• Metals: good conductors of heat and electricity  

• Nonmetals: bad conductors of heat and electricity  

• Metalloid: in between metal and nonmetal properties (they are okay  conductors of heat and electricity).

• Group 1A: the alkali metals (Li, Na, K, Rb, Cs, and Fr) • Group 2A: the alkaline earth metals (Be, Mg, Ca, Sr, Ba, and Ra) • Group 6A: the chalcogens (O, S, Se, Te, and Po)

• Group 7A: the halogens (F, Cl, Br, I, and At)

• Group 8A: the noble gases (He, Ne, Ar, Kr, Xe, and Rn) • Groups 1B, 3B-8B: transition elements/transition metals • Mole: unit of measurement that shows the quantity in any substance  

that has the same number of particles (atoms) found in 12.000 grams  of carbon 12. It is like a dozen (12) or a gross (144).  

• Avogadro’s number: the number of atoms in 12.000 grams of carbon  12 (6.022 x 10!!)

• Molar mass: the mass in grams of one mole of a substance. • Kinetic energy: results from motion  

• Potential energy: is the energy an object has at a still position  • Chemical energy: energy which is stored in the structural units of  chemical substances

• Electrostatic energy: energy which results from the interaction of  charged particles

• Law of conservation of energy: the total amount of energy in the  universe is constant, and it cannot be created or destroyed. • Joule (J): the SI unit for energy, which symbolizes the amount of  kinetic energy that is possessed by a 2 kg mass moving at 1 m/s • Electromagnetic spectrum: a continuum of radiation, which includes  radio waves, microwaves radiation, infrared and ultraviolet radiation.  • Wavelength (λ): the distance between two peaks of a wave  • Frequency (ν): the number of waves that pass through a certain point  in 1 second.

• Amplitude: the vertical distance from the middle of the wave to the  top of the peak  

• Electromagnetic wave: a wave which possess an electric field and  magnetic field component

• Blackbody radiation: the radiation a solid gives off when it is heated

• Photoelectric effect: a theory which states that when electrons are  exposed to light on a metal surface, they are ejected from that  surface

• Threshold frequency: the minimum of light that can cause electron emission  

• Photons: little particles of light  

• Emission spectra: the heat given off a piece of sample material when  it is energized using either thermal or other form of energy • Line spectrum: the emission of light only at certain wavelengths  • Ground state: as an electron gets closer to the nucleus it becomes  more negative, and at its most stable energy point

• Excited state: the energy state, which is more positive. The electron  jumps from the ground state to the excited state.

CHAPTER 1-3.4 IMPORTANT FORMULAS AND NUMBERS

Base Quantity

Name of Unit

Symbol

Length

meter

m

Mass

kilogram

kg

Time

second

s

Electric current

ampere

A

Temperature

kelvin

K

Amount of  

substance

mole

mol

Luminous  

intensity

candela

cd

Prefix

Symbol

Meaning

Tera-

T

1×10!"

Giga-

G

1×10!

Mega-

M

1×10!

Kilo-

k

1×10!

Deci-

d

1×10!!

Centi-

c

1×10!!

Milli-

m

1×10!!

Micro-

μ

1×10!!

Nano-

n

1×10!!

Pico-

p

1×10!!"

• To convert from Celsius to Kelvin use this formula: K = C + 273.15 • To convert Celsius to Fahrenheit, use the formula: !!× ℃ + 32 = ℉ • To convert Fahrenheit to Celsius, use the formula: !!× ℉ − 32 = ℃ • To convert Fahrenheit to Kelvin, use the formula:

K = 59× ℉ − 32 + 273.15

• Density is the ratio of mass to volume, you calculate density using  the following formula: d = !!

• The mass of an electron is 9.10 × 10!!"g

• An atomic radius is about 100 pm in radius

• The radius of the atom’s nucleus is 5 × 10!!pm

• The highest known density of an element is 22.6 g/cm!, which is for  iridium.  

• You calculate average atomic mass (deals with isotopes) by this  formula:

((normal mass) x (natural abundance of normal x 100)) + ((isotope  mass) x (natural abundance of isotope x 100))

• Avogadro’s number: (6.022 x 10!!)

• Kinetic energy formula: Formula: E! = !!mu! 

• Electrostatic energy: E!" α !!!! 

!

• 1 Joule= !"×!! 

!! 

• The speed of light (c) is equal to 2.99792458×10! m/s • Speed, wavelength, and wave frequency are related by this formula:  c = λν

• The energy of a single quantum of energy is show by this formula:  E=hν

• Energy binding electrons in a metal, that has enough energy to  knock the electron lose can be shown by this formula: hv = KE + W • The Rydberg equation: !! = R!( !!!! - !!!!)

• The energies that the electron in the hydrogen atom contains can be  shown by this formula: E! = −2.18×10!!"J ( !!!)

• Look at equations 3.8 and 3.9 in book! (Page 73)

***I did not use sig figs!

***I did not use sig figs!

Lecture / Book Notes: Chapter 2 (1/25/2016)

CHEM 1030 Cagg

Highlighted: Vocab ----- Highlighted: People ----- Highlighted: Formula/NumbersThe first test WILL NOT heavily emphasize sections 2.1-2.4, so you can skim through  them. It WILL heavily emphasize 2.5-2.7!!!

Section 2.1 

❖ Atoms

• Atoms: are the building blocks and smallest units of matter

- The philosopher, Democritus was the first to propose this idea

- An English scientist and school teacher, John Dalton, was the first to  agree/formalize with the idea of matter having atoms  

- Atoms are made up of tiny subatomic particles

• Element: contain a unique form of the same atom (each element has a  different form of atoms), which cannot be divided into 2 or more simpler  substances

- Example: oxygen, nitrogen, iron, sulfur

Section 2.2 

❖ Discovery of electrons  

• Radiation: is the transmission of energy in the form of waves  

• Cathode ray tube: also known as the electron beam, is a glass tube with  metal plates on each end, which has all air, sucked out from it. This device  detects electrons using radiation.  

• Cathode: is a negatively charged plate within the cathode ray which  produces radioactive waves called cathode rays  

• Anode: is a positively charged plate within the cathode ray  

• Columb’s law: like charges push against one another and opposite charges  come towards one another  

- Example: magnets  

• The rays in a cathode ray tube, are a stream of negative particles (electrons),  which was discovered by the English physicist J.J. Thomson 

- Thompson determined the charge-to-mass ratio of the electrons as  1.76 × 10! Columbs per gram 

• R.A. Millikan who was an American physicist, determined the charge of an  electron by examining the motion of oil drops

• The mass of an electron is 9.10 × 10!!"g 

❖ Radioactivity  

• Wilhelm Rontgen, a German physicist discovered X-rays.  

• X-rays: rays that are able to penetrate matter and cause materials to give off  a fluoresce light. They do no have charged particles due to them not being  deflected by anything.

- X-rays are only produced when exposed to cathode rays

• Antoine Becquerel, a French physicist discovered radiation  

• Radioactivity: a spontaneous emission of rays that are highly energetic and  cannot be deflected due to anything. There are 3 types of radioactive  emissions:

- Alpha α rays: have positively charged particles, which are called alpha  particles. These are deflected away from a positively charged plate. - Beta β rays: have negatively charged particles called beta particles (also  known as electrons). These are deflected away from negatively charged  plates.

- Gamma y rays: these have no charge

❖ Proton and the nuclear model of the atom

• The plum pudding model: an idea proposed by Thomson saying that  electrons were embedded into the atom like “raisins in a scoop of rum ice  cream”; it was the accepted theory for quite a while.

• Ernest Rutherford, a student of Thomson’s and a physicist, claimed that  atoms were mostly empty space, he also said the atom had a very dense  core called the nucleus, with positive particles (protons) in it.

• An atomic radius is about 100 pm in radius

• The radius of the atom’s nucleus is 5 × 10!!pm 

• Protons and neutrons are located INSIDE of the nucleus, and the electrons are located AROUND the nucleus.

❖ The neutron

• James Chadwick, an English physicist discovered the neutrons (which are  neutral particles) in the nucleus  

Section 2.3 

❖ Atomic number, mass number, and isotopes

• The atomic number: the number that defines the element, it also represents  the number of protons that are in the nucleus of that atom.

- Example: Carbon’s atomic number is 6, so a carbon atom has 6 protons in  its nucleus

- Note that the number of protons can NEVER change, or it will change the  element, so if you added one more proton to carbon, it would then have  7 protons and become nitrogen.

• The atomic mass number: the total number of protons and neutrons that are  in the nucleus of an element’s atom.

• The nucleus usually contains has the same amount of protons as it does  neutrons, so carbon has 6 protons, and 6 neutrons; however, this is not the  case with isotopes.

• Isotopes: are atoms which have the same atomic number, but different  atomic mass, (the number of neutrons is different from the number of  protons).

• To find the protons of an element, look at the atomic number

• To find the neutrons of an element, subtract the atomic number from the  atomic mass.

• To find the electron of an element, look at the atomic number, there is the  same amount of protons as there are neutrons.

Section 2.4 

❖ Nuclear Stability

• The nucleus has most of the atom’s mass, but it’s a small part of an atom’s  total volume.

• The highest known density of an element is 22.6 g/cm!, which is for iridium. • The stability of an atom’s nucleus is determined by the difference between  coulombic repulsion and short range attraction.

Section 2.5 

❖ Average atomic mass

• Atomic mass: the mass of an atom in atomic mass units (amu)

• You calculate average atomic mass (deals with isotopes) by this formula: ((normal mass) x (natural abundance of normal x 100)) + ((isotope mass) x  (natural abundance of isotope x 100)) 

- To calculate the average atomic mass of carbon 12 and carbon 13 you  would do:

((12.00000 amu) x (98.93 x 100)) + ((13.003355 amu) x (1.107 x 100)) = 12.01  amu

Section 2.6 

❖ The periodic table  

• The periodic table: consists of 118 elements that are grouped according to  their physical and chemical properties.

• The elements are arranged into periods (horizontal) and groups/families  (vertical).

• Metals: good conductors of heat and electricity  

• Nonmetals: bad conductors of heat and electricity  

• Metalloid: in between metal and nonmetal properties (they are okay  conductors of heat and electricity).

• Group 1A: the alkali metals (Li, Na, K, Rb, Cs, and Fr)

• Group 2A: the alkaline earth metals (Be, Mg, Ca, Sr, Ba, and Ra) • Group 6A: the chalcogens (O, S, Se, Te, and Po)

• Group 7A: the halogens (F, Cl, Br, I, and At)

• Group 8A: the noble gases (He, Ne, Ar, Kr, Xe, and Rn)

• Groups 1B, 3B-8B: transition elements/transition metals

Section 2.7 

❖ The mole  

• Mole: unit of measurement that shows the quantity in any substance that has  the same number of particles (atoms) found in 12.000 grams of carbon 12. It is like a dozen (12) or a gross (144).

• Avogadro’s number: the number of atoms in 12.000 grams of carbon 12  (6.022 x 10!!) 

❖ Molar mass

• Molar mass: the mass in grams of one mole of a substance.

• Molar mass and Avogadro’s number can be used to convert to and from  mass, moles, and number of atoms.

Lecture Notes: Chapter 1 (1/13/2016)

CHEM 1030 Cagg

Highlighted: Vocab ----- Highlighted: Rare Usage ----- Highlighted: Formula

Section 1.1  

❖ The scientific method  

• A series of 6 steps used to approach an experiment  

1. Observation  

2. Hypothesis  

3. Experiment  

4. Develop a theory  

5. Further experiment  

Section 1.2  

❖ States of matter

• There are three states of matter that a substance can be

- Solid: particles tightly packed with no space in between (ice)

- Liquid: particles are packed loosely with some space in between (water  at room temperature)

- Gas: particles are not packed and roam freely (water vapor)

• We can convert a substance from one state of matter to another without  changing the identity of the substance.

- When you melt an ice cube, the state of matter changes from solid to  liquid, but it doesn’t change the original substance, which is water.  ❖ Mixtures  

• Homogenous mixtures: these mixtures are thoroughly mixed to where you  cannot tell the original substances (also called a solution).

- Example: sugar and water

• Heterogeneous mixtures: these are not thoroughly mixed, and you can often  tell what the original substances were.

- Example: sand and water

Section 1.3  

❖ Properties of matter

• Properties of a substance can be either quantitative or qualitative - Quantitative properties: are measured or expressed using a number,  ALWAYS ADD UNITS!  

o Example: you have a handful of almonds, to determine the  

quantitative property of them, you would count to see how many  

there are.

- Qualitative properties: are based mainly on observation  

o Example: you have a basket of different colored apples, to  

determine the qualitative property of them; you would look at the  

colors, the sizes, etc.

• Physical properties  

- Physical properties: are those, which can be observed/measured  without changing the identity of the substance. Some types of physical  properties are: color, melting point, boiling point, smell, etc.

o Example: you can determine the freezing point of water by  

taking water in its liquid state and then freezing it  

- Physical changes: are those which disrupt the state of matter of a  substance (the substance PHYSICALLY changes).

o Example: when you melt an ice cube, it goes from a solid rock  

to a puddle of water

• Chemical properties  

- Chemical properties: is a reaction that a substance has when  

interacting with another substance.

o Example: gasoline is flammable, so when you bring it close to  

fire, the fire will intensify  

- Chemical changes: reactions that result in the total change of the  composition of the substance, so that the original one does not exist  anymore.

o Example: when food is being digested, acids come to break the  food down, and the food is changed to a completely different  

substance than it once came in.

• Extensive and intensive properties

- Extensive properties: depend on how much matter there is in an  object, the more matter an object has the more its mass is  

o Two objects of the same mass (extensive property) can be  

added together, mass and volumes are 2 types

▪ Example: you have 2 apples, you add up the masses and  

in result get sum of the individual masses of each apple  

- Intensive properties: don’t depend on how much matter there is in an  object, temperature and density are 2 types

o If you have two graduated cylinders filed with distilled water both  at the same temperature, and you combine them into a large  

beaker, the density and temperature WILL NOT change, and  

stay the same as it was when the two were separate.

Section 1.4  

❖ SI base and prefix units  

• “The International System of Units” is the revised metric system  • There are 7 base SI units

Base Quantity

Name of Unit

Symbol

Length

meter

m

Mass

kilogram

kg

Time

second

s

Electric current

ampere

A

Temperature

kelvin

K

Amount of  

substance

mole

mol

Luminous intensity

candela

cd

Highlighted: rarely if not ever used it class (said by professor)

• To measure length use a meter stick  

• To measure volume use a biuret, graduated cylinder, pipet, or volume flask  • To measure mass use a balance

• To measure temperature use a thermometer

• There are 10 prefixes used with SI units  

Prefix

Symbol

Meaning

Tera-

T

1×10!"

Giga-

G

1×10!

Mega-

M

1×10!

Kilo-

k

1×10!

Deci-

d

1×10!!

Centi-

c

1×10!!

Milli-

m

1×10!!

Micro-

μ

1×10!!

Nano-

n

1×10!!

Pico-

p

1×10!!"

Highlighted: rarely if not ever used it class (said by professor)

❖ Mass: the measure of the amount of matter in an object  

• The mass of an object stays the same no matter where it is - Example: a TV remote weighs 2 kg on Earth, it will weigh 2 kg on  Mars.

• Weight and mass are NOT the same thing  

- Weight: the force which is exerted by an object due to gravity, the  weight of an object varies from place to place since gravity is involved

o Example: a dog weighing 5 kg on Earth may weigh 10 kg on  

Mars

• 1 kg = 1000g = 1×10! g

• The most commonly used unit is not an SI unit, it is an amu (atomic mass  unit), which is equal to 1.6605378×10!!"g or 1.6605378×10!!"kg

❖ Temperature

• There are 2 main temperature scales used in chemistry, Celsius and Kelvin  - The C and K scales are equal in magnitude, so if the temperature of  tea increases by 5 degrees Celsius, then it will also increase by 5  

Kelvin. Note that C and K are NOT the same!!!

- The boiling point of water on the Celsius scale is 100 degrees C,  freezing point of water on the Celsius scale is 0 degrees C.

- Absolute value on the Kelvin scale -273.15

- Kelvin CANNOT be negative

- To convert from Celsius to Kelvin use this formula: K = C + 273.15 • Fahrenheit is another temperature scale, but not used as much in chemistry - Fahrenheit is the most common unit for temperature used in the US - Boiling point of water on Fahrenheit scale is 212 degrees F, freezing  point of water on Fahrenheit scale is 32 degrees F

- To convert Celsius to Fahrenheit, use the formula: !!× ℃ + 32 = ℉ - To convert Fahrenheit to Celsius, use the formula: !!× ℉ − 32 = ℃ - To convert Fahrenheit to Kelvin, use the formula:

K = 59× ℉ − 32 + 273.15

❖ Volume and density  

• Volume and density require base units that are not in the base SI units, so we  have to derive units by combining base units.

- Derived SI unit for volume is m! 

o Liter (L) is also used, this is derived from dm! 

o Milliliter (mL) is also used, this is derived from cm! 

• Density is the ratio of mass to volume, you calculate density using the  following formula: d = !!

- Example of density: what happens when you mix oil with water? The  oil floats on top because it has a lower density

- Derived SI unit for density is kg/m!, but this unit is a little large so we  often use g/cm! 

Section 1.5  

❖ Significant figures

• Significant figures: are the MEANINGFUL digits in a number

• There are many rules to obtaining the correct amount of significant  digits/figures in a number

- When you don’t have a zero in the number, all digits are significant o Example: #1 13457.98 (7 SF)

o Example #2: 1789.3 (5 SF)

- When zeros are placed in between non zero numbers, then they are  counted as significant  

o Example #1: 103.2 (4 SF)

o Example #2: 10205 (5 SF)

- Zeros on the left of a non zero number are not counted as significant  o Example #1: 0.031 (2 SF)

o Example #2: 0.1 (1 SF)

- Zeros on the right of the non zero number are counted as significant IF  the number has decimal points

o Example #1: 1.450 (4 SF)

o Example #2: 1.0 (2 SF)

- Zeros on the right of the last non zero number which does not contain  a decimal point, MAY or MAY NOT be counted as significant  

o Example #1: 5000 (can have 1, 2, 3, or 4 SF, since no additional  information is given, we cannot tell)

o To avoid complications of this and for significant figures to  count, write the number in the following ways

▪ If you are showing 1 SF, write the number as 5 x 10! 

▪ If you are showing 2 SF, write the number as 5.0 x 10! 

▪ If you are showing 3 SF, write the number as 5.00 x 10! 

▪ If you are showing 4 SF, write the number as 5.000 x 10! 

- When adding or subtracting  

o The answer can only have as many decimal places as there are  in the original number with the lowest amount of numbers on the  right of the decimal place  

▪ Example #1: 4.32 + 8.236 = 12.556 (to show the correct  

SF, you will have to round your answer to 12.56,  

because from the original number, the lowest amount of  

numbers after the decimal are 2, from 4.32)

▪ Example #2: 5.6777 - 2.1 = 3.5777 (to show the correct  

SF, you will have to round your answer to 3.6, because  

from the original number, the lowest amount of numbers  

after the decimal place is 1)

- When multiplying or dividing

o The answer can only have as many SF as there are in the  

original number with the lowest amount of SF

▪ Example #1: 6.1 x 8.93 = 54.473 (to show the correct SF,  

you will have to round your answer to 54, because the  

lowest amount of significant figures in your original  

number is 2)

▪ Example #2: 10.22225 / 4.345 = 2.35264672 (to show the  

correct SF, you will have to round your answer to 2.353,  

because the lowest amount of significant figures in your  

original number is 4)

- When dealing with exact numbers

o Exact numbers have an infinite number when dealing with  

significant figures. Exact numbers are ones that don’t contain  

decimal places like 3, 102, 57, 500, 65

▪ Example #1: If you have 6 apples that have a mass of  

10.3 g each, how much is the total mass of the 6 apples?  

6 x 10.3 g = 61.8 g (you can report this number as 62, or  

61.80, or 61.800, SF amount does not matter because 6  

is an exact number)

▪ Example #2: 7 x 51.237 = 358.659 (you can report this  

number as 359, 358. 66 or, 358.6590, SF amount does  

not matter because 7 is an exact number)

❖ Accuracy and precision  

• Accuracy: shows how close or exact a measurement is to the real value  - Example #1: you have a bowl of candy contains 60 pieces, when you  count the candy in the bowl, you also count 60 pieces, so your answer  is accurate.

- Example #2: think of a bulls eye board, if you throw 3 darts, and they  all hit the center, then you have throw very accurate shots

• Precision: shows how close a set of measurements are to one another  - Example #1: you have a set of data containing the numbers 12.34,  12.35, 12.34, 12.35, 12.36, the data is very precise due to how close  the numbers are to one another.

- Example #2: think of a bulls eye board, if you throw 3 darts, and they  hit almost the center, but not quite, then you have thrown very precise shots.

 

 

 

Section 1.6  

❖ Conversion factors

• Conversion factors: fractions where 2 quantities that mean the same thing  (equal to each other) are put one in the numerator and other in the  denominator.

- Example #1: 12 inches = 1 foot ???? !" !" 

! !"

- Example #2: 1 meter = 10 decimeters ???? ! ! 

- Example #3: 1 pound = 16 ounces ???? ! !" 

!" !"

!" !"

❖ Dimensional analysis

• Dimensional analysis: also known as the factor label method, is a way of  converting units in problem solving  

• Make sure to use SF when doing this!

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