Chemistry Reviewer (finals 1).docx

GENERAL CHEMISTRY I FINALS REVIEWER | © DANE DE MATA FINAL EXAM COVERAGE - Electron configuration ends with either ns1 or 1. The Periodic Table of Elements ns2 2. Parts of the Periodic Table - n = period # = energy level 3. Periodic Trends - Group 1 & 2 4. Noble Gas Electron Configuration P BLOCK 5. Valence Electrons & Octet Rule - The p orbital can hold a maximum of 6 6. Lewis Dot Diagrams electrons 7. Ionic Bonds → 6 groups in the p block 8. Covalent Bonds - Electron configuration ends with np1 until np6 9. VSEPR Theory - n = period # = energy level 10. Electron and Molecular Geometry - Group 13-18 11. and Hybridization D BLOCK 12. Molecular and Bond Polarity - The d orbital can hold a maximum of 10 electrons THE PERIODIC TABLE OF ELEMENTS → 10 groups in the p block - Electron configuration ends with (n-1)d1 until Periods - Horizontal rows (n-1)d10 - Seven periods - n = period # - Period # indicates the energy level - Outer transition metals - PERIODICITY: the repeating characteristics F BLOCK of elements - The f orbital can hold a maximum of 14 electrons Groups/Families - Vertical columns → 14 groups in the p block - Group # indicates the no. of electrons in the - Electron configuration ends with (n-2)f1 until outermost shell (n-2)f14 ● GROUP 1A & 7A through 0 - n = period # ○ Representative elements - Inner transition metals ○ Readily form compounds due to their number of valence electrons ● GROUP B ○ Transitions metals ○ Inner and outer ○ Quite rare; compounds containing them are limited Dmitri Mendeleev (1834-1907) - Listed the elements in several vertical columns in order of increasing ATOMIC MASS Father of the Periodic Table - Immortalized in element 101 Henry Moseley (1887-1915) - Determined the nuclear charge of elements - Arranged the elements in a table by order of ATOMIC NUMBER - Conceived the structure of the modern periodic table Periodic Law - When the elements are arranged in order of increasing atomic number, there is a periodic pattern in their physical and chemical properties PARTS OF THE PERIODIC TABLE Blocks 1. S BLOCK - The s orbital can hold a maximum of 2 electrons → 2 groups in the s block Page 1

GENERAL CHEMISTRY I FINALS REVIEWER | © DANE DE MATA Representative Elements GROUP 1 - Alkali metals - Very reactive - Ionic charge of +1 GROUP 2 - Alkaline Earth metals - Less reactive than Group 1 - Ionic charge of +2 GROUP 18 - Noble gases - Inert or non-reactive - All shells are completely filled - 8 valence electrons - Xenon and Krypton have compounds, but they are unstable GROUP 17 - Halogens - Very reactive - Ionic charge of -1 GROUP 16 - Chalcogens - Oxygen family - Less reactive than Group 17 - Charge is -2 GROUP 15 - N family; +3, -5 GROUP 14 - C family; +4, -4 GROUP 13 - B family; -3, +5 *METALLOIDS - B, Si, Ge, As, Sb, Te, Po - Has properties of both metals and nonmetals Transition Elements INNER TRANSITION ELEMENTS - Generally radioactive - Composed of the Lanthanides and Actinides OUTER TRANSITION ELEMENTS - Most are hard with high densities - Often magnetic - High melting and boiling points PERIODIC TRENDS - Trends exist in some periodic properties - Physical and chemical behavior based on the electron configurations - electron configurations are used to explain many of the repeating periodic properties of the elements - PERIODIC LAW: properties of element are a function of their atomic number (also their electron configurations) Atomic Radii - Radius of an atom cannot be measured directly - Atom doesn’t have a defined boundary because the electrons are in a cloud

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DOWNWARD ○ Higher energy levels of valence electrons ○ Higher energy level, electrons are FARTHER from the nucleus ● LEFT to RIGHT ○ Outer shell is at the same level, but as you go right there are more protons added Ionic Size - CATION (+) ANION (-) - Ions are formed by losing or gaining electrons - Metallic elements readily form cations (LEFT SIDE) - Non-metallic form anions (RIGHT SIDE) - Cations are smaller than the parent atom ● The pull of the protons increases (divided between less electrons) - Anions are bigger than the parent atom - ISOELECTRONIC SERIES ● The more positive an ion, the smaller it will be ● Conversely, the more negative, the larger it will be Ionization Energy - Energy required to remove an electron from an element (to make it positive or more positive) - First ionization energy: IE1 is the energy to remove the FIRST ELECTRON ● Elements with small IE1 tend to form cations; and those with large IE1 form anions ● Opposite with atomic radii Example: Na or K → Na has a higher ionization energy

*EXCEPTIONS - Be (s2) to B (s2 p1) because s shields p ● Outer electrons can easily escape (low IE) ● Trend goes the opposite Page 2

GENERAL CHEMISTRY I FINALS REVIEWER | © DANE DE MATA - N (s2 p3) to O (s2 p4) because repulsions exist between the first paired electrons OCTET RULE ● There are breaks in the trend of - Octa = 8 ionization energy - Elements aim to form a noble gas electron configuration in order to achieve stability - This is true for representative elements Electronegativity - Ability of an element to attract electrons when - Transition metals are prone to have multiple bonded with another atom charges - Arbitrary scale: highest score is 4.0 for - s + p orbitals: COMPLETELY FILLED Fluorine EQUALS 8 - ENx < ENY *EXCEPTIONS X-Y - The ff elements: He, H, Li, Be, B want to form a 1s2 configuration - INCOMPLETE octet and SUPER octet (more on these later)

NOBLE GAS ELECTRON CONFIGURATION - Shorter notation that uses Group 18 - Useful for when the electron configuration is long He

2

1s2

Ne

10

2p6

Ar

18

3p6

Kr

36

4p6

Xe

54

5p6

Rn

80

6p6

-

After writing the noble gas in brackets, start with the succeeding s orbital - For writing the NGEC a noble gases, use the preceding one Examples: - N (7) EC: Is2 2s2 2p3 NGEC: [He] 2s2 2p3 - Hg (80): [Xe] 6s2 4f14 5d10 - Ar (18): [Ne] 3s2 3p6 VALENCE ELECTRONS - Outermost electrons - Electrons on the s and p sublevels - Group # gives the number of valence electrons

LEWIS DOT DIAGRAMS - Dots represent the electrons - X = added electrons - Convention: if there is a charge, use brackets and label the charge on the outside (upper right)

IONIC BONDS - Electron TRANSFER between a metal and nonmetal - Electronegativity difference (ED) determines the type of chemical bond > 1.7 = IONIC < 1.7 = COVALENT - Ignore the # of atoms when looking for ED NaCl

2.1

Ionic

H2O

1.4

Covalent

BeCl2

1.5

Covalent

Ionic Compounds - Ionic bond is found between the ions (electrostatic attraction) - Not between the transfer of electrons

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GENERAL CHEMISTRY I FINALS REVIEWER | © DANE DE MATA - Never call them molecules → FORMULA UNITS - NaCl → Sodium Chloride

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CATIONS: no need to draw the valence electrons because the outer shell already either has O or 8 ANIONS: still draw the electrons, including the ones that are added Examples: POTASSIUM SULFIDE

MAGNESIUM OXIDE

Criss Cross Method

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CRYSTAL LATTICE STRUCTURE ● Network of ions bound together - Name of the compound is the CATION + ANION (ide) - When writing the formula of ionic compounds, do NOT include the charges Examples: K2O → Potassium Oxide Al2S3 → Aluminum Sulfide Na3N → Sodium Nitride Na2(CO3) → Sodium Carbonate Stock System - For metalloids with multiple charges, roman numerals represent the charge - Initially, the ratio might be in lowest terms so CHECK THE ANION to determine if the charge is correct Examples: IRON (II) CHLORIDE: FeCl2 → Fe+2 Cl-1 CHROMIUM (II) OXIDE: CrO→ Cr+2 O+2

SODIUM CHLORIDE

-

-

C+m + A-n Combine to form the compound CnAm The ratio of n:m must be in lowest terms in order to ensure the # of atoms are correct Group

Fixed Charge

IA

+1

IIA

+2

IIIA

+3

IVA

+4, -4

VA

Polyatomic Ions - Contain compounds, usually anions - Such as oxides of nitrogen, sulfur and phosphorous N

S

P

Suffix (ite)

NO2-1

SO3-2

PO3-3

(Suffix ate)

NO3-1

SO4-2

PO4-3

-

Also including oxyhalides (chlorate, borate, iodate) with the formula where n is 1-4 XOn-1

ClO-1

hypochlorite

-3

ClO2-1

chlorite

VIA

-2

ClO3-1

chlorate

VIIA

-1

ClO4-1

perchlorate

Naming & Formula Writing

Nomenclature for Acids IONIC BINARY ACIDS Page 4

GENERAL CHEMISTRY I FINALS REVIEWER | © DANE DE MATA - Similar to ionic binary compounds, but the COVALENT BONDS metal is replaced by Hydrogen (H) - Involves the SHARING of valence electrons between two nonmetals - HYDRO + ANION (ic acid) Examples: - Sharing is done simultaneously within the Cl-1 (anion name: chloride) electron cloud → hydrochloric acid - A “transfer” is happening a million times per Br-1 (anion name: bromide) second → hydrobromic acid - Each atom contributes one electron to form a shared pair IONIC TERNARY ACID ENDING WITH SUFFIX - 0 < ED < 0.4 = NONPOLAR COV IDE - 0.5 < ED < 1.7 = POLAR COV - HYDRO + POLYATOMIC ION - Drop the -ide, replace with -ic acid Examples: CN-1 (anion name: cyanide) → hydrocyanic acid IONIC TERNARY COMPOUNDS ENDING WITH SUFFIX -ITE - POLYATOMIC ION - Drop the -ite, replace with -ous acid Examples: ClO-1 (anion name: hypochlorite) → hypochlorous acid NO2-1 (anion name: nitrite) → nitrous acid IONIC TERNARY COMPOUNDS ENDING WITH SUFFIX -ATE - POLYATOMIC ION - Drop the -ate, replace with -ic acid Examples: ClO3-1 (anion name: chlorate) → chloric acid ClO4-1 (anion name: perchlorate) → perchloric acid Ionic Compound Properties 1) High melting point ● Due to the crystal lattice structure ● Difficult to vibrate all the atoms 2) Dissolves easily ● Aqueous or molten but does not melt ● Good electrical conductor ● Can carry the charge in a circuit since the atoms can now move freely within the solution 3) High brittleness ● Ions of different charges next to each other separate easily

Types of Bonds - # of shared pairs is equal to the # of bonds ● SINGLE BOND ● DOUBLE BOND ● TRIPLE BOND - There is a maximum of 3 shared pairs or 6 electrons - Single bonds are LONGER than double and triple bonds - The triple bond is the SHORTEST BOND which is the least stable or most reactive - When drawing, one line = one bond = two electrons - Attempts to follow the OCTET RULE to achieve a more stable state

Lewis Structures Group

Bonding Pairs

Non-bonding Pairs

C

4

0

N

3

1

O

2

2

F

1

3

H

1

0

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GENERAL CHEMISTRY I FINALS REVIEWER | © DANE DE MATA *Examples:

*Examples

Exceptions to the Octet Rule 1) Hydrogen is an exception to the octet rule since it tries to mimic Helium (only two valence electrons) to achieve stability 2) Incomplete Octet ● Less than 8 electrons ● Result is that it is very reactive ● For example, B can only pair 3 electrons even if its needs 2 more 3) Super Octet ● More than 8 electrons ● Possible due to the d orbital Naming & Formula Writing - Name is the (PREFIX) CENTRAL ATOM + (PREFIX) BONDED ATOM (-ide) - The prefix mono- is not included for the central atom Number

Greek Prefix

1

Mono-

2

Di-

3

Tri-

4

Tetra-

5

Penta-

6

Hexa-

7

Hepta-

8

Octa-

9

Nona-

10

Deca-

Common Name

IUPAC Name

H2O

water

Dihydrogen monoxide

CH4

methane

Carbon tetrahydride

N2O

Nitrous oxide

Dinitrogen monoxide

NH3

ammonia

Nitrogen trihydride

N2H4

hydrazine

Dinitrogen tetrahydride

Covalent Compound Properties 1) Low melting and boiling point ● Due to relatively weak intermolecular forces ● Covalent compounds are usually found at room temperature in 3 states of matter 2) Inability to conduct electricity ● There are no charged particles that are capable of transporting electrons VSEPR THEORY - Valence Shell Electron Pair Repulsion - Outermost electrons consist of bonding and nonbonding or lone pairs - The ELECTRON REPULSION of lone pairs is greater than that of bonding pairs - Thus, atoms within a molecule will form a specific geometry that MINIMIZES the repulsion between them - They can achieve this by positioning themselves as far from each other as possible - Take note that this is within THREE DIMENSIONS ELECTRON & MOLECULAR GEOMETRY Electron Geometry - Generalized shape of the compound - A → central atom X → # of ELECTRON DOMAINS - Electron domains include bonds and lone pairs AX2-6 Molecular Geometry - Actual shape of the compound, taking into consideration REPULSIVE FORCES - A → central atom B → bonds E → lone pairs AB2-6E0-3 Types AX2

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GENERAL CHEMISTRY I FINALS REVIEWER | © DANE DE MATA Mol. Geom AB2

Shape Linear

Angle

AB2E3

Linear

180

180

AX3 Mol. Geom

Shape

Angle

AB3

Trigonal Planar

120

AB2E

Bent/V-shape

<120

AX6 Mol. Geom

Shape

Angle

AB6

Octahedral

90

AB5E

Square Pyramidal

<90

AB4E2

Square Planar

90

AX4 Mol. Geom

Shape

Angle

AB4

Tetrahedral

109.5

AB3E

Trigonal Pyramidal

<109.5

AB2E2

Bent/V-shape

<109.5

Types of Bonds SIGMA BOND (σ) - Represents single bonds in the compound - Head to head overlap of electron clouds PI BOND (π) - Represents double and triple bonds - Side to side overlap of reserved electrons

AX5 Mol. Geom

HYBRIDIZATION - Hybridization explains a compound’s molecular geometry based on its electron configuration (valence)

Shape

Angle

AB5

Trigonal Bipyramidal

90, 120

AB4E

See-Saw

<90, <120

AB3E2

T-Shape

<90

Hybridized Orbitals Electron Domains

Hybridization of the Central Atom

Electron Geom

2

sp

AX2

3

sp2

AX3 Page 7

GENERAL CHEMISTRY I FINALS REVIEWER | © DANE DE MATA 4

sp3

AX4

5

sp3d

AX5

6

sp3d2

AX6

Steps for Hybridization 1) Draw the Lewis Structure of the molecule and count the # of electron domains of the CENTRAL ATOM 2) Based on the # of electron domains, determine the HYBRIDIZATION (ref. to table above) 3) Draw the ORBITAL DIAGRAM of the valence electrons of the central atom 4) Encircle the orbitals to be hybridized and reserve electrons as needed. Remember: ● S subshell → 1 orbital ● P subshell → 3 orbitals ● D subshell → 5 orbitals 5) Draw the new HYBRIDIZED ORBITAL with properly distributed electrons. Filled orbitals indicate the # of lone pairs, while unpaired electrons represent the sigma bonds. *Note: In step 4, reserve electrons based on the # of π bonds present in the compound Example: Hybridization of an sp3 orbital

Electron Cloud Notation - Shapes are similar to electron geometry - This notation shows the head to head and side to side overlaps that occur during hybridization

Example: Hybridization of Water

Step-by-Step 1) Formula is H2O 2) Oxygen is the central atom in the Lewis structure and has 4 electron domains.

3) Thus, its hybridization is sp3. 4) The orbital diagram of oxygen’s valence electrons is:

5) When hybridized, this becomes:

which shows that there are 2 lone pairs and 2 electrons available for bonding in its molecular geometry.

MOLECULAR & BOND POLARITY

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GENERAL CHEMISTRY I FINALS REVIEWER | © DANE DE MATA Bond polarity is determined by the electronegativity difference (ED) of the various atoms

0 < ED < 0.4 = NONPOLAR COV 0.5 < ED < 1.7 = POLAR COV 1.7 < ED = IONIC -

Molecular polarity is based on the VSEPR Theory and a molecule’s geometry

Non-Polar Molecules - Even distribution of electrons - Atoms have no partial charge - The net dipole moment is equal to ZERO. Polar Molecules - Electrons are attracted to a particular atom - Uneven distribution of electrons - Net dipole moment is NOT ZERO. What is a DIPOLE MOMENT? - It is the VECTOR that shows the tendency of electrons to move towards the more electronegative atom. - Notation is an arrow with a plus sign at the end. Examples:

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For a molecule to be non-polar, the sum of all these vectors MUST EQUAL ZERO. - Dipole moments cancel out in 2 situations: 1) Two vectors in the opposite direction are exactly 180° apart 2) All vectors are of equal distance from each other. Examples: BF3 CO2

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