CHEM 115 Spring 2014 Lecture # 1

Class Announcements

• Make sure you read the syllabus in its entirety…you are responsible for knowing the information in it

• Be sure to check out the Chem 115 eCampus website regularly: • https://ecampus.wvu.edu• Class notes, answer keys, grades, etc. are posted on this website

• Class notes and handouts for this week are posted!• Date class notes!!

• Register for the ALEKs online homework:• An instructional handout is posted on eCampus in the “ALEKS”

folder• Course Codes given in instructional handout on eCampus

and in email• Complete Objective 0 by January 12th at 11:59PM

Laboratory

Chem 115 Lab DOES NOT meet the first week of classes….Actual Lab Starts the week of January 15th or 17th

1. Your not in lab this week BUT…complete Appendices A and B at home (in the back of your lab manual). • Due at the BEGINNING of lab (January 15th or 17th )• Work does not need to be shown.

2. Lab room and desk numbers assignments will be posted on eCampus sometime next week…so you know where to go for lab!

3. Complete Pre-Lab Assessment # 0– Read instructions (posted on eCampus)

4. Study for Lab Quiz – Questions will be similar to the Pre-Lab Assessment questions and will include questions from the syllabus!

Purchase lab goggles, apron, and lab manual from the UNIVERSITY BOOKSTORE:• White aprons from Book Exchange are not permitted• A binder might be a good idea too!

Additional Announcements

• No attendance will be taken the first week of classes but you should bring your ID and a non-programmable calculator to lecture and lab from now on!

• Get them out at the beginning of each lecture

• Chemistry Learning Center (CLC): MTW 7-10 pm FREE HELP!

• Will let you know when this starts

• PLTL: www.pltl.wvu.edu

• Will let you know when this starts

• CHEM Tutors for Hire: Clark Hall Room 217

• I can help you pick a tutor if you would like too!

Chapter 1

Types of Numbers Used in Chemistry

Exact Numbers:

• A number with a value that is exactly known

• No error or uncertainty in the value

• Numbers obtained by counting individual objects AND/OR defined numbers within a given measurement system

Measured Numbers

• Number with a value that is NOT exactly known due to the measuring process.

• Some error or uncertainty in the value

• Amount of error depends on the measuring device (increment and distance between markings) and you!

Significant Figures, and Uncertainty

• Measured numbers will ALWAYS contain some uncertainty or error

• This error is expressed by the number of significant figures (or digits) reported for the measured number

• Significant Figures: Digits used to represent a measured number such that only the digit farthest to the right is uncertain1. The digit farthest to the right is an estimate, which we assume to

have an error of plus or minus one (±1)

2. The digit farthest to the right is significant in some instances and insignificant in some instances…I will teach you!

3. As the number of significant figures increase, uncertainty/error decreases, and the precision of the measuring device increases

Rules for Determining Sig Figs

• Rule for Determining the Number of Significant Figures :

1. All nonzero digits are significant

• Example: 4.13 3.222 15124875

2. Different Types of Zeros:

• Trailing Zeros: 4.130 3.2220000 15.00

• Captive Zeros: 4.022 3006 15.890201

• Leading Zeros: 0.082 0.000222 0.890201

• Zeros with Decimal: 550.00 44000. 32.0

• Ambiguous Zeros (No Decimal): 20 5500 330022000

- Measured numbers greater than one that end in zero(s) without a decimal

place are ambiguous and MUST be written in standard/scientific

exponential form: 20 5500 2005 101 5500.0

2.0 X 101 5.500 X 103

Rounding

1. If the number that is removed is 5 or greater increase the last retained digit by 1 (round up)• Example: 1.3667 (4 SF) =

2. If the number that is removed is less than 5 leave the last retained digit unchanged (round down)– Example: 1.3664 wants 4 SF =

ROUND AT THE VERY END OF THE CALCULATION!!

These rules are different than the textbook!

Significant Numbers and Calculations

• Addition/Subtraction Calculations:

• Rule: The final answer has the same number of digits past the decimal place as there are in the measurement with the fewest digits after the decimal place

• Example: 33.2 + 66.23 = ? What do you do?

1.Do the calculation

• Answer = 99.43

2.Find the measurement (from the original numbers) with the fewest digits after the decimal place

Make sure you ALWAYS set up your addition/subtraction problems properly by lining up the decimal places!!!!!!!!

Significant Figures and Calculations

• Multiplication and Division Calculations:

• Rule: The final answer contains the same number of SIGNIFICANT FIGURES as there are in the measurement with the fewest number of significant figures

• Example: Volume (cm3) = 9.2 cm X 6.83 cm X 0.3774 cm = ?

1. Do the calculation

• Answer = 23.7143064 cm3

2. Find the measurement (from the original numbers) with the fewest significant figures

Multistep Calculations

• When performing multistep calculations do addition/subtraction FIRST followed by multiplication/division

• DO NOT ROUND UNTIL THE VERY END!

• Example:

(6.9371 + 0.30) X 0.01689 =

• Example:

0.7761 X 22.1 =

49.01 – 48.89

Standard Scientific Exponential Notation

• Standard Exponential Form (Scientific Notation): # x 10n

• # is the coefficient• Coefficients are not always written properly; however, you

should follow the rules given

• n is the exponent

• Converting numbers greater than 1 into scientific notation:

1. Move decimal place to the left until you have a number between 1-10 (This is your coefficient)

2. Add the times you have moved the decimal place as n (This is your exponent)

• #’s > 1 have positive exponents

• Example: 233 =

Scientific Exponential Notation

• Converting numbers less than 1 into scientific notation:

1. Move the decimal place to the right until you have a number between 1-10 (This is your coefficient)

2. Subtract the number of times you have moved the decimal place as n (This is your exponent)

• #’s < 1 have negative exponents

• Example: 0.233 =

• There are some exceptions and you should pay attention to them!!

Examples

• Convert the following into scientific notation or to a measured #:• 0.002536• 0.1234• 0.000000006298• 12578• 100• 89000654000• 9.678 X 10-3

• 3.45 X 102

• 0.0000012556 X 104

Calculations with Scientific Notation

• Multiplication of Numbers in Scientific Notation:

1. Multiply the coefficients

2. Add the exponents

• Division of Numbers in Scientific Notation:

1. Divide the coefficients

2. Subtraction the exponents

4.0 X 104 X 2.0 X 105 =

6.0 X 106 /2.0 X 103 =

Calculations with Scientific Notation

• Addition/Subtraction of Numbers in Scientific Notation:

1. Convert the numbers so they have the same exponent• Convert the smaller exponent to the larger exponent

2. Add/subtract the coefficients (Remember SF rules!)

3. Keep the exponent the same

• Multistep Calculations of Numbers in Scientific Notation:

1. Do addition/subtraction FIRST followed by multiplication/division…round at the end (if necessary)!

0.23 X 104 + 4.5 X 105 =

Examples

• Determine the number of sig figs in the following:

1. 900 2. 0.0030 3. 0.1044 4. 53069

5. 0.00004715 6. 0.0000000071600 7. 30.070

8. 200 9. 101 10. 18.933202001

• Perform the following calculations using rounding rules:

1. 1.04 X 0.5588 = 2. 13.13557 / 101.30 =

3. 1.0438 – 0.5588 = 4. 6.1 X 104 + 2.16 X 103

3.615 X 102

Some More Numbers

• The Seven Fundamental SI Units of Measure

• The International System of Units:

• Based on the metric system, which is convenient for most people

• We use these units so that all chemists can “speak” to each other

• Recognize these units of measure!

Measurements

• You can use prefixes to account for very large or small #’s:

• You don’t always obtain a measurement in SI units and so…. you have to convert the measurement to the desired unit!

• These prefixes can act as CONVERSION FACTORS in dimensional analysis!

Measurements: Unit Conversion or Dimensional Analysis

• If a conversion factor is correctly applied, a unit will always cancel

Conversion Factor: 1 mile = 5280 feet• 1760. feet = ? miles

• 0.453 mile = ? feet

• Here is an example where the units are NOT correctly applied:

• 1760. feet = ? miles

Measurements: Unit Conversion that Require Multiple Steps

• Several conversion steps (and this will usually be the case) may be necessary to solve a problem:

• How many seconds are there in 7 days?

• How many mm are there in 2.5 km? (You need to know these conversion factors!)

Measurements: Temperature

• Temperature (T): A measure of how hot or cold one object is relative to another• A thermometer is the most common means for measuring

temperature• Temperature is NOT the same thing as heat (q), which is the energy

that flows from the object with the higher temperature to the colder object with the lower temperature

• How do they work?• When immersed in a substance hotter than the thermometer, heat

flows from the substance through the glass into the fluid in the thermometer causing the fluid to expand and rise

• Opposite occurs when immersed in a substance colder than the thermometer

Measurements: ThreeTemperature Scales

Measurements: Converting Temperature

• Three temperature scales: Celsius, Kelvin, and Fahrenheit• Based on the physical state of water• Celsius (ºC)

• Boiling point: 100 ºC• Freezing point: 0 ºC

• Kelvin (K)• Absolute zero: 0 K• Boiling point: 373.15 K• Freezing point: 273.15 K

• All measurements are positive values and have no º sign• Uses the same size degree unit as the Celsius scale

• Conversion from K to ºC or ºC to K:• K = ºC + 273.15• ºC = K – 273.15

Measurements: Converting Temperature

• Fahrenheit (ºF) – Differs from Celsius and Kelvin scale in its zero point and in the size of the unit of measurement

• Boiling point: 212 ºF

• Freezing point: 32 ºF

• When converting from Fahrenheit to Celsius or Kelvin, you must adjust for the start point

• 180/100 = 9/5

• 100/180 = 5/9

• Conversion from ºF to ºC or ºC to ºF:

• ºF = (ºC X 9/5) + 32

• ºC = (ºF – 32) X 5/9

Measurements: Temperature Example Problems

• Example: A child has a body temperature of 38.7 ºC. If normal body temperature is 98.6 ºF, does the child have a fever?

• Convert ºC to K

Chemistry: Properties of Matter

Chemistry is the study of matter and its changes

• Physical Properties: Characteristics that do not involve a change in a sample’s chemical makeup

• Melting Point Electrical Conductivity Solubility

• Density Temperature Hardness Odor

• Physical Change: Occurs when a substance alters its physical form (rearranges its molecules), NOT its composition


• Chemical Properties: Characteristics that do involve a change in a sample’s chemical makeup

• Rusting Combustion Tarnishing

• Chemical Change: Any change that results in the formation of new chemical substances• Chemical changes involve making or breaking bonds between

atoms

• Example: Balloons and Traffic Light Demonstration


• Intensive Properties: Independent of sample size

• Temperature Color

• Melting/boiling point Hardness

• Extensive Properties: Dependent on sample size

• Length

• Volume

• Height

Chemistry: Matter and Its Three Physical States

1. Solid - Has a fixed shape that does not conform to the container shape (definite volume and shape)

• The molecules are very close together (touching) and cannot move around; they are tightly packed (can’t flow)

• Not compressible

2. Liquid - Conforms to the container shape but fills the container only to the extent of the liquids own volume (definitive volume but no definite shape)

• The molecules are close together (could be touching) and they move around (can flow)

• Not extremely compressible

3. Gas - Conforms to the container shape, and it fills the entire container (no definite volume or shape)

• Molecules are widely separated and can move around freely • Compressible

Chemistry: States of Matter

Chemistry: States of Matter

• Physical States of Water:

• Chili

Measurements: Density

• Density (d) – The density of an object is its mass divided by its volume

Density is commonly measured in g/cm3 or g/mL1 cm3 = 1 mL

• You can use the knowledge of density to your advantage:• A substance whose density is greater than that of the liquid

will sink in the liquid• A substance whose density is less than that of the liquid will

float on the liquid.• A substance whose density is equal to that of the liquid will

remain wherever placed in the liquid and will neither sink nor float.

Measurements: Density

• You can use mass and volume to determine density:

• You can use mass and density to determine volume:

• You can use volume and density to determine mass:

Chapter 2

Classification of Matter

Matter

Pure Substances Mixtures

HomogeneousMixtures

HeterogeneousMixtures

Compounds Elements

What Are Elements Really Made Of?

Greek philosophers like Plato, Aristotle, and Democritus speculated about what an element was made of… but technology was an issue in 400 B.C.

Much later it was found that….

Elements are composed of atoms

• Atomic Theory provides us with ideas about the structure, properties, and behavior of atoms

We Know What an Atom is Made Of ! (i.e. Structure)

• An atom is electrically neutral and consists of three subatomic particles

• Different atoms have different amounts of subatomic

particles!

• In a neutral atom, # protons = # electrons

because the overall charge of an atom = 0

Charge Mass (amu)

Mass (grams) Location

Proton +1 1.00728 1.67 X 10-24 Nucleus

Neutron 0 1.00867 1.67262 X 10-24 Nucleus

Electron -1 0.0005486 9.10939 X 10-28 Outside nucleus

The Periodic Table and Atomic Notation

• Atomic Number (Z) : the number of protons in the nucleus of an atom

• Example: Boron =

Sodium =

Cobalt =

• Mass Number/Atomic Mass Number (A): the total number of protons and neutrons in the nucleus of an atom

• Example: Boron =

Sodium =

Cobalt =

• Determining # of Neutrons:

# Neutrons = Mass # (A) – Atomic # (Z)

Atomic symbol

SyAZ

You can determine the number of each subatomic particle for every atom using the periodic table

So how do you calculate the number of electrons???

What are Isotopes?

• Isotope: atoms of an element having:• The same number of protons • Different number of neutrons

So….• The SAME atomic number (Z) • DIFFERENT mass number (A)

• Determining # of Neutrons: • # Neutrons = Mass # (A) – Atomic # (Z)

• Example: Carbon has three naturally occurring isotopes12C, 13C, 14C

The number of neutrons in isotopes are different; however, the protons and electrons remain the same!

Atomic Mass Unit (amu)

If we consider the actual mass of the proton, neutron, and even electron we would have rather large atomic masses for the elements so….

Atomic masses of all elements were referenced to the atomic mass of the most abundant isotope of carbon (12C) in amu

for convenience• 6 protons and 6 neutrons

Atomic Mass Unit (amu): Based on Carbon-12 or 12C

1 atom 12C = 12 amu (exactly)OR

1 amu = 1/12 the mass of an atom of 12C = 1.660539 X 10-24 g

• Example: If the relative mass of Mo:12C is 7.995, what is the atomic mass of Mo on the 12C atomic mass scale?

Average Atomic Mass Number (A)

Why is the atomic mass of carbon given as 12.011 amu instead of 12 amu?

Mass Number (Atomic Mass # (A) ): The weighted average of the isotopic masses of the element’s naturally occurring isotopes

• Atomic masses shown on periodic table are average atomic masses taking into account the different isotopes of each element and their percent abundances (isotopic abundances)

Examples

• Calculating Average Atomic Masses:• Example: It is found that carbon consists of two naturally occurring

isotopes (12C and 13C) with atomic masses and % abundances given below. Calculate the average atomic mass of carbon

Isotope Atomic Mass % Abundance12C 12 amu 98.89%13C 13.0034 amu 1.11%

Examples

• Calculating % Abundance:• Example: A sample of naturally occurring gallium has an average

atomic mass of 69.7 and consists of two isotopes, gallium-69 and gallium-71. Given the information shown below, calculate the % isotopic abundances of the two isotopes

Isotope Atomic Mass69Ga 68.9 amu71Ga 70.9 amu

Studying the Atom… But How Did We “Find” the Subatomic Particles?

• J. J. Thomson used cathode-ray tubes• When a magnet or an electrically charged plate is around the tube the

cathode ray is deflected• This means that because the cathode ray is produced at the negative

electrode (cathode) and deflected to the positive electrode (anode) there must be negatively charged subatomic particles in the ray, which we now call electrons

Atoms contain negatively charged electrons

The charge to mass ratio (e/m) = 1.758820 X 108 C/g

The Atom: Electron

Thompson could measure charge to mass ratio NOT the actual charge or mass of an electron but R. A. Millikan did

• Millikan was able to calculate the mass of each droplet and use the known voltage applied to the plates to determine the charge of an electron. Using Thompson’s charge to mass ratio he then determined the mass

Charge of an electron = -1.60 X 1019 C

Mass of an electron = 9.11 X 10-28 g

The Atom: Protons and Neutrons

If an atom is electrically neutral but electrons are negatively charged you have to have a positive charge somewhere to balance the electrons charge

• Rutherford discovered the nucleus using alpha particles (+ charge) which he shot at gold foil. Most went through the gold foil undeflected but some (1 of every 20,000) were deflected back

• As the alpha particles went toward the gold foil they were deflected because they encountered the gold’s positively charged nucleus and were repelled (like charges repel)

Most of the atom is empty spaceIn center of atom there is a massive, positively charged core called the nucleus

The Size of the Nucleus


Matter


HomogeneousMixtures


Compounds Elements

Matter: Elements & the Periodic Table

• Element is the simplest type of matter with unique physical and chemical properties that consist of only one type of atom

• Elements cannot be broken down

• The periodic table organizes the elements and are represented by symbols

Organization of the Periodic Table

• Groups are vertical columns• Periods are horizontal row

• Metals: Shiny solids at room temperature (Hg is exception)• Conduct heat and electricity• Approximately ¾ of the periodic table are metals• Does not include H

• Non-Metals: Gases or dull brittle solids at room temperature (Br is exception)

• Conduct heat and electricity poorly• Includes H

• Metalloids: Have properties between those of a metal and non-metal• Also called semimetals• B, Si, Ge, As, Sb, Te, Po, At



• Alkali Metals: Metals in Group IA (Li, Na, K, Rb, Cs)• Have similar properties

• Alkaline Earth Metals: Metals in Group IIA (Be, Mg, Ca, Sr, Ba)• Have similar properties:

Both Alkali Metals and Alkaline Earth Metals are found in the S-Block of the periodic Table

• D-block: Transition Metals

• F-block: Lanthanide and actinide Metals


Alkali Metals


Alkaline Earth Metals


• Halogens: Non-metals in Group VIIA (F, Cl, Br, I, At)

• Have similar properties

• Exist as diatomic molecules in elemental form at RT (F2, Cl2, Br2, I2)

• Noble Gases: Non-metals in Group VIIIA (He, Ne, Ar, Kr, Xe, Rn)

• Have similar properties: not very reactive and do not readily form compounds

• Exist as monatomic gases at RT (He(g), Ne(g), Ar(g), etc.)

• Both Halogens and Noble Gases are found in the P-Block of the periodic Table

• The P-block contains metals, semi-metals and non-metals


Halogens


Noble Gases


Matter


HomogeneousMixtures


Compounds Elements

Matter: Mixtures

• Mixture is a group of two or more pure substances (elements and/or compounds) that are PHYSICALLY intermingled

• No chemical changes of the individual substances

• NOT chemically combined

• NOT in combined in fixed proportions by

mass

• Types of Mixtures:

1. Homogenous mixtures

2. Heterogenous mixtures


Matter


HomogeneousMixtures


Compounds Elements

Matter: Compounds

• Compound is a substance that is composed of two or more elements in fixed proportions that are CHEMICALLY combined

• Compounds can be broken down through chemical changes

• The chemical and physical properties of compounds are different from the elements that formed them

Compounds: Chemical Formulas and Formula Units

• Chemical Formula – Shorthand way of writing the chemical symbols of each element in a compound

• Subscripts are used to indicate the number of atoms of each element present in the compound

• If no subscript is given, the number 1 is understood

• Also called a formula units for ionic compounds

• Also called a molecular formula for molecular compounds

• Example – H2O NaCl Ca(OH)2

Compounds: Ionic vs Molecular

Depending on what type of bond is formed will determine the type of compound formed

• Ionic Compounds: An electrostatic attraction between a positive ion and a negative ion, where one or more electrons have been transferred from the valence shell of one atom to the valence shell of the other atom• An ionic compound typically forms between a metal and non-

metal• Characteristics: good electrical conductors if the ions are mobile (liquids,

solutions). Tend to be hard, brittle, often with high melting points

• Molecular Compounds: Sharing valence electrons between atoms of different elements form COVALENT bonds• Typically formed between NON-METALS

• Characteristics: poor electrical conductors; small molecules may be gas, liquid or soft solids

Ionic Compounds: Cations and Anions

What is an ion and how are they formed?

Ionic Compounds: Cations and Anions

• Ion: A charged species that forms when an atom gains or loses an electron • The ion becomes charged either positively or negatively

• Cation - positively charged ion • Formed when atoms lose electron(s)• Metals tend to form cations

• Anion - negatively charged ion• Formed when atoms gain an electron(s)• Non-metals tend to form anions

Example: How many of each type of subatomic particle

(# e-, # p, and # n) are present in the following?

• 16O-2 32S-2 137Cs+1

Ionic Compounds: Why Do Atoms Gain or Lose Electrons?

• Every element wants to obtain noble gas configuration and an “octet” WHY?

• Because NGC/octet is the most stable and lowest in energy• In order to obtain this, elements will donate or accept electrons

(valence electrons) from the valence shell

Ionic Compounds: How Many Electrons are Gained or Lost?

• Monatomic Ions: Ions consisting of ONE atom

• Example - Cl-, Al+3, Zn+2, N-3

Charges on monatomic ions can be predicted from positions in the periodic table… so you know how many electrons were lost or

gained!!!

Metals NonmetalsIA IIA B-Groups IIIA IVA VA VIA VIIA VIIIA

+1 +2 Variable +3 -4/+4 -3 -2 -1 0

(Fe+2/Fe+3)

Transition and post-transition metals (metals in Groups IIIA –VIA) have variable charges

Example - Sn+2/Sn+4, Sb+3/Sb+5

Usually forms covalent compounds

Ionic Compounds: Charges of Ions

+1 +2 +3 -4/+4 -3 -2 -1

Some transition and post-transition metals form more than one cationSome elements would rather SHARE than TRANSFER their electrons!

Ionic Compounds: Polyatomic Ions

• Polyatomic Ions: Ions consisting of two or more elements• Atoms are held together by molecular/covalent bonds BUT form

ionic compounds! • So when asked what types of bonds are present in a compound with a

polyatomic ion…you say both!

• The charge shared over all atoms in the ion and atoms stay together as a unit

Example - NO3-, NO2

-, CO3-2, PO4

-3, SO4-2, SO3

-2, etc.

• Brackets are used when 2 or more polyatomic ions appear in the ionic formula unit

• Ex. Ca(NO3)2 = 1 Ca2+ cation & 2 NO3- polyatomic anions

• Sorry, but you are going to have memorize the list of common polyatomic ions • You need to know the name, formula, and charge• A incomplete list can be found on last page of the syllabus or pg. 62 in

the textbook

Ionic Compounds: Common Polyatomic Ions You Need to Know

Ionic Compounds: Transferring Electrons and Neutral Compounds

In the formation of sodium chloride, one electron is transferred from the sodium atom to the chlorine atom BUT…. the overall charge of an ionic compound = 0!!

Na1+ + Cl1-Na + Cl

11 protons10 electrons




Molecular Compounds: Covalent Bonds

Molecular compounds (aka Covalent compounds) have covalent bonds

• Covalent Bond: A form of chemical bonding characterized by sharing of valence electrons between atoms

• Typically formed between NON-METALS

• Molecule is the smallest bit of a molecular

compound

Molecular Compounds: Diatomic Molecules

• All of the gaseous nonmetallic elements except for the noble gases exist as diatomic (2 atom) molecules

• Know these!

Compounds: Example Problems

• Identify whether the following are molecular/covalent or ionic compounds:

1. KCl

2. H2O

3. CaCl24. FeSO4

5. CO2

6. F2

7. CH4

8. NH3

9. AgNO3

Compounds: Naming Ionic Compounds

• In the name of an ionic compound, the cation is always given first and the anion second

1. The cation is always specified as the name of the metal and is NOT changed

2. The anion is specified by using the first part of the nonmetal name and then adding the suffix “ide”• EXCEPTION: Don’t change ending of polyatomics ions!!

3. Examples: NaCl Al2O3 MgS

For ionic compounds, you do not indicate the numbers of each type of atom…. So NO prefixes!

Compounds: Naming Ionic Compounds

• Transition metals can often have several stable positive ions (variable charge)… that is why we skip the D-block when assigning charge!

• Example - Fe2+ and Fe3+, Cu1+ and Cu2+

1. The modern system for naming indicates the charge on the cation as a roman numeral within the name of the substance:

• Iron (II) chloride FeCl2 Copper (I) chloride CuCl

• Iron (III) chloride FeCl3 Copper (II) chloride CuCl2

2. An older system, still in common use, uses the suffixes –ous and –ic to indicate the charge on the cation:

• Ferrous chloride FeCl2 Cuprous chloride CuCl

• Ferric chloride FeCl3 Cupric chloride CuCl2

Balancing Chemical Formulas

• When balancing the chemical formulas for ionic compounds, you must consider charges…this is different with covalent compounds

• The overall charge of the compound must be = to zero!

• Use the criss-cross method for generating balance chemical ionic formulas

• Example:

You can also used a balanced chemical formula to determine the charge of a transition element!

Compounds: Naming Molecular Compounds

Naming Molecular Compounds:

1. Name the element that comes first in the chemical with prefix

– Mono is never used to name the 1st element

2. Name 2nd element 2nd but add -ide to end of name and use prefix

3. Example: N2O4

CS2

S3F4

Cl2O6

1 Mono

2 Di-

3 Tri-

4 Tetra-

5 Penta-

6 Hexa-

7 Hepta-

8 Octa-

9 Nona-

10 Deca-

Prefixes

Some Special Covalent Compounds: Acids

• Acid – A substance that provides H+ ions in water (aq)

• Nomenclature (Naming) of Acids (aq):

• Hint: H usually written first in chemical formula of acids

1. Decide whether acid is a BINARY or OXYACID

Binary Acid – contains H and one other element.

Ex. HCl(aq), HF(aq), H2S(aq)

Oxyacid – contains H, O, and one other element.

Ex. HClO4, H2SO3, HNO2

• FYI:

• Base – A substance that provides OH- ions

in water (aq)

Acids Cont’d

Naming of binary acids:

Use root of nonmetal name.

Add hydro- prefix and –ic suffix to root followed by acid

Examples: Provide names for the following binary acids

HCl(aq) HF(aq) HBr(aq)

Acids Cont’d

Naming of oxyacids:

1. Use root of nonmetal name followed by acid

2. Usually four different oxyacids can be formed from nonmetal, specify which is present by adding following prefixes and suffixes.• hypo- -ous oxyacid with least oxygens• -ous • -ic• per- -ic oxyacid with most oxygens

Example: Provide names for the following oxyacidsHClO2, HClO4, HClO3, HClO, H2CO3 and H2SO3

• Most polyatomic ions are anions derived from oxyacids by removal of one or more protons (H+)

“ate”“ic”

“ite”“ous”

Compounds: Naming Polyatomic Ions

Something To Think About In Lab…

In the Pursuit of Studying the Atom…We Found Out About Compounds Too

• Fundamental Chemical Principles:

• Law of Mass Conservation – Mass is neither created nor destroyed in chemical reactions

• Antoine Lavoiser

• Example: 71.8 g iron oxide decomposes to 55.8 g iron and 16.0 g oxygen

• Law of Definite Proportions – Different samples of a pure chemical substance always contain the same proportion of elements by mass

• Joseph Proust

• Example: Carbon dioxide contains: 27.3 % carbon

72.7 % oxygen

How Can We Explain These Laws?

• Dalton proposed a theory of matter to explain the Law of Mass Conservation and the Law of Definite Proportions

He came up with:1. Elements are made up of tiny particles called atoms that cannot be created

or destroyed

2. Each element is characterized by the mass of its atoms • Atoms of the same element have the same mass, but atoms of different

elements have different masses• Example: All carbon atoms have a mass of ~12.0107 but oxygen

atoms have a mass of ~15.9994

3. Atoms of one element cannot be converted into atoms of another element

4. Chemical combination of elements to make different substances occurs when atoms join together in small whole-number ratios• 2 hydrogen atoms and 1 oxygen molecules form water

Another Law: The Law of Multiple Proportions

• Elements can combine in different ways to form different substances, whose mass ratios are small whole-number multiples of each other

• Example: Benzene and ethane combine carbon and hydrogen in different ways

• Benzene combines 4.61 g of carbon and 0.39 g of hydrogen• Mass ratio of C/H = 4.61 g carbon/ 0.39 g hydrogen = 12 g C/H (2 SF!!)

• Ethane combines 4.00 g of carbon and 1.00 g of hydrogen• Mass ratio of C/H = 4.00 g carbon/1.00 g hydrogen = 4.00 g C/H (3 SF!!)

• Law of Multiple Proportions:• Mass ratio of C/H in benzene : ethane = 12 g C/H = 3 X as many carbon

4.00 g C/H in benzene This would only make sense if it is true that matter is composed of discrete

atoms that have characteristic masses and combine with each other in specific ways

Dalton’s Atomic Theory explains known observations AND can accurately predict the outcome of experiment not yet conducted….This is what we need

in a Law!

Documents

CHEM 115 Spring 2014 Lecture # 1