Chemistry 1301a

Midterm 2 flashcards - chapter 2.2 (periodic trends) and chapter 3 (structure and bonding)

what are the five periodic properties

1. effective nuclear charge

2. atomic/ionic size (bond lengths)

3. ionization energy

4. electron affinity

5. electronegativity

Dmitri Mendeleev and the Periodic Table

• whats he called

• whatd he based it on

• coined the father of the PI

• based it off of density and reactivity

what is the periodic law

when the elements are arranged in order of increasing atomic number (Z) certain sets of properties recur periodically

modern PI arrangements

• what are the rows

• what are the vertical rows

+what mass does the mass number show

• periods= rows of elements arranged by increasing atomic number (Z)

• groups- the vertical rows

mass number shown is the more stable artificial isotope of the element

what are main group elements

elements in the s and p block of the PI

what groups are in s block

• Group 1: alkali metals (one valence e-)

• Group 2: alkaline earth metals (two valence e-)

what groups are in p block

• group 13: the triels (3 ve)

• group 14: the tetrels (4 ve)

• group 15: the pnictogens (5 ve) …”to choke”

• group 16: the chalcogens (6ve)…”ore forming”

• group 17: the halogens (7 ve)…”salt forming”

• group 18: the noble gases (8 ve/full valence shell)

what blocks are the transition metals in

+what orbitals to the ve tend to occupy

block d and block f

d block: transition metals

f block: inner transition metals

ve tend to occupy the d and f orbitals

two types of e in many electron atoms

• valence e and core e

valence: occupy outer orbitals, further from the nucleus

core (Ncore) : occupy inner orbiatals, close to nucleus

explain

effective nuclear charge

• net positive charge experienced by valence e shielded by core e from full charge of nuc

• Z* = Z - Ncore , where Z is the number of electrons in total and Ncore is all the inner electrons (not in valence shell)

• as we move from left to right on periodic table, core e Ncore = same ; nuc charge Z increases and therefor Z* also increases

• higher Z* value, higher attraction b/w valence e and atomic core

define

size of atoms

atomic radius

size of ions

1. atomic radii

2. half the length of a bond formed by 2 atoms of the same element

3. charctreriszed by ionic radii; ionic radius is approx from crytal lattice structure parameters of ionic cmp

relationship between atomic size and Z*

as you move from left to right on the period table, radii decreases. as Z* increases, nuc attraction b/w core + valence e increases

• valence e are pulled closer ot nuc..,overall atomic size decreases

why is the Ga atom smaller than Al despite being before it on the period table?

• since d and f orbitals have poor shielding properties, p electrons outside of d-electron shells are not shielded as much as expected adn are strongly attracted to the nucleus

• there for smaller atomic radii and smaller atom

characteristics of single, double and triple bonds (+atom size and bond length)

• larger atom = longer bond length

• double and triple bonds tend to be smaller than single bonds becuase they contain more electrons that pull nuclei together
  
  

define

anion

cation

anion

• e added

• need more space for these e so always larger than neutral parent

cation

• e lost

• need less space so always smaller than neutral parent

• same number of + but less e so more pull on the e that are still there

explain

size of isoelectronic species

• as Z increases, species decrease in size

• because pull on e from nucleus containing fixed number of e increased with Z… e cloud decreases in size

generalizations for atom size with di/mono an/cat-ions

for the same element

• dications< monoanions

• dianions > monoanions

for isoelectronic species

• the more negative species is the largest

• the most posiive speciies is the smallest

define

ionization energy

+consider PT and the properties as you travel along it

• energy required to remove valence e from a gas-phase atom ( or ion)

• second IE required more energy than first IE (bc ur taking a secon electron after already taking one…remaining e are held closer to nucleus)

• takes less e to steal from a larger atom becuase their e are super far from the nucleus

• since it is related to how e interact with attraction to nucleus, as you go left and right on the PT, IE increases

• as you go down PT, IE decreases… as atoms get larger the distance between valence electronds and the nucleus increases so it is easier to overcome the electron-nucleus attraction and remove a valence e)

anomalies in IE

• boron has lower IE than Be…it has a 2p electron which occupies a higher energy level than the 2s electrons…easier to ionize

• ionizing O is easier than N…neutral oxygen has one p orbital that contrains 2 electrons. This creases strong e repulsion in that orbital. the p orbials are half filled which is favorable…so removing the electron in O is easier than N

define

electron affinity

• energy (enthalpy) change when gaseous atom gains an electron

• enthalpy (ΔH)

exothermic (negative ΔH) is favorable

endothermic (positive ΔH) is not favorable

• the E required to add eacg successive e is less favorable

e want o be as far away as possible and you keep adding more which is bound to make more in the same orbital…more energy

• usually increases over a period w the exception of noble gases and anomalies, decreases down a group

define

electroneagativity

• ability of an atom to draw a pair of e of a bond to itself

• ability of an atom participating in a chemical bond to draw electron density to itself

• helps explain reactivity

• pauling electronegativities (most popular way to classify)

• smaller atoms attract e more strongly than larger atoms, electronegativity tends to decrease with increasing atomic size

• increases in rows and decreases down groups

which e participate in bonding

only the valence e

connection between shape and properties of molecules

• chemical, physicial and biological proerties of all molecules depend on their shape ad electronic structure

• molecular shape is determined primarily by the spacial arrangment of the atomic nuclei

• electronic structure of a molecule is determined mainly by the interactions among the valence electrons of constituent atoms

explain how

electronigativity values determine the types of bonds formed

• elements of low EN (alkali and lalkaline earth metals) typically form ionic bonds

• elements of moderate EN (carbon, nitrogen, phosphorus) typically form covalent bonds

• elements of higher EN (oxygen, halogens) can form covalent or ionic bonds

define

ionic bonds

• complete transfer of e from one atom to another

• occurs when EN difference between atoms is greater than about 1.9

define

crystal lattice

+magnitude of force b/w anions and cations formula

• formed by the strong electrostatic formes of ionic bondsng

• solid state and 3D

• high melting point for ionic compounds

magnitude of force b/w anions and cations formula

• coulombs law

• ionic compounds that involve smaller ions tend to have larger attraction forces and therefore higher melting points

relationship between force b/w molecules and melting/boiling point

• larger force = stronger attraction between ions = higher melting point

• higher melting point means more energy is required to distrupt attraction (melt)

define

covalent bonds

+types of covalent bonds

• EN difference is less than about 1.0

• e shared between bonded atoms

• polar covant and nonpolar covalent

explain

the difference (characteristics) between polar and nonpolar covalent bonds

nonpolar covalent

• equal sharing of e density

• EN ≤ 0.5

polar covalent

• unequal sharing of e density

• 0.5 ≤ EN ≤ 1.9

• forms a dipole moment (e goes towards the atom with higher EN)

characteristics of a dipole moment

• molecules that have it will orient themselves in an electric field

• dipoel moment has both direction and magnitude so its treated as a vector … drawn as an arrow

define lewis structures

• diagram repersentations of chemical structures

• NOT intedned to convery shape info

• they do show wher ethe bonds are, how many bonds, if there are lone pairs and also formal charge

explain

octet rule

• if possible, atoms should have (be able to claim) a filled valence shell of e

• 2 e for the first row elements

• 8 e (octet) for the secon row elements

• second row cannot have moe than 8 e (due to their orbitals)

• 3rd row or greater can have more than an octet (elements have additional energetically accessible orbitals that can accomodate additional electrons (expanded octet))

define

formal charge

• e accounting cocepy (not necessarily a real charge)

• differnece between the number of e in an atoms free, neutral state and the number of e that “belong” to it in the Lewis strucutre

• Cf = N valence - N owned

• where N owned is the total number of nonbonding electrons that resude on that atom plus half the number of electrons in all the covalent bonds formed by the atom

method to determining formal charge

1. count the total number of bonds formed by the atom. a single bond counts as one, a double bond counts as 2 and a triple bond counts as three

2. count the number of nonbonding electrons. all nonbonding electrons bleond to the atom they reside on

3. the sum on 1 and 2 is N owned, the total number of electrons owned by the atom

4. the number N valence of ana atom is equal to the second digit of the group number of the chemical element

rules to formal chage

1. all nonbonding e belong to atom

2. half of the bonding e belond to atom (e in covlaent bond are shared)

3. check that the sun of formal charges equal overall charge

rules to drawing lewis structures

1. Sum the total number of valence electrons (including charges)

2. Draw skeleton structure (single bonds)

• Usually central atom is C or atom with lowest electronegativity

3. Determine how many electrons remain unused

4. Distribute remaining electrons

• Add non-bonding electrons to give full octet (except for H), starting with most electronegative atoms

• Place remaining electrons on central atom

• Make double or triple bonds as needed to meet octet rule

• remember that elements in the third row are allowed to have more than 8 electrons

5. Assign formal charges to atoms (non-zero)

• Make double or triple bonds to best distribute formal charge

• Check that formal charge sum matches overall charge

deifne

expanded octet

• expanded octect are permitted to make the best lewis strcutre

• only applies to elements in the third period

explain how

to choose the best lewis structure

1. full octect on all atoms (first priority)

2. best distribution of formal charge

• least internal formal chage seperation

• most EN atom gets negative formal charge

• otherwise structure will be energitically unfavaorable

  

  
  
  

define resonance structures

• you can draw multiple correct lewis structures that differ ONLY in the position of the electrons

• electrons are waves so the e can move around and both structures are valid
  
  

define

resonance hybrid

• charge is delocalized over both atoms, bonds are identical and on bonding e are not shown

• 
  

  

  
  
  

explain

bond order

Single bond = 1

Double bond = 2

Triple bond = 3

Average bond order of one bond type (in a compound)

= (sum of (specific) bond orders) / (number of connections)

• for resonance hybrids, since it is a composite of structures, the term average bond order is used to desbrice a simple average number of bonds bw specific atoms

rules to identify and read line diagrams

• each solid line is a covalent bond

• if no atom is drawn at the intersection of two lines, it is assumed to be a carbon atom

• if no atom is drawn at the end of a line, it is assumed to be a carbon atom

• hidrogen atoms bonded to carbons are not shown (there are usually enough H to fill the four bonds of C)

• all non-carbon atoms (heteroatoms) must be drain along with hydrogens that are attached to them

• nonbonding electrons are often ommited, but are assumed to be there. the number of nonbonding electrons must be consistent with the formal charge of the atoms

why do we have VSPER

-what does it stand for

• Shape is predicted by VSEPR theory, which examines the pairs of electrons in the valence shell of a central atom (O, in the case of H2O)

• VSEPR: Valence Shell Electron Pair Repulsion

what is VSEPR theory

1. pairs of electrons, both bonding and non bonding (lone pairs), occupy the valence shell of a central atom

2. these electron pairs, also called regions of electron density, repel each other and specially positioj themsleves to maximize seperation of the pairs

3. lone pairs exert greater repulsive forces than bonding pairs

4. multiple bonds get treated as a single region of electron density

explain

AXE notation

• AX(m)E(n)

• A denotes the central atom

• n is the total number of loe electrons pairs, E, on the central atom

• m is the number of atoms, X bonded to central atom A

• the sum of m+n is the total numbrer of regions of electron density around the central atom

explain

difference between molecular shape and electron geometry

+how to determine molecular shape

Molecular shape MAY be different from electron geometry (if central atom has lone pairs of electrons)

To determine molecular shape

• Draw the Lewis Structure (bonding/non-bonding regions)

• Determine electronic arrangement (# regions of electron density)

• Determine molecular shape (bonding regions of electron density)

explain

2 regions of e density

m + n = 2

• Electronic arrangement: Linear (180°)

• Molecular shape and AXE options: AX (2) E

• sp hybridization

explain

3 regions of electron density

m + n = 3

• Electronic arrangement: trigonal planar (120°)

• sp2 hybridization

  

  ]

• trigonal planar

• bent (1 lone pair)

explain

4 regions of electron density

m + n = 4

• Electronic arrangement: tetrahedral (109.5°)

• sp3 hybridization

• tetrahedral

• trigonal pyramidal ( 1 lone pair)

• bent (2 lone pairs)

explain

5 regions of electron density

m + n = 5

• Electronic arrangement: trigonal bipyramidal (90° and 120°)

• sp3d hybridization

• trigonal bypyramidal

• seesaw (1 lone pair)

• t shape (2 lone pairs)

• linear (3 lone pairs)

explain

six regions of electron density

m + n = 6

• Electronic arrangement: octahedral (90°)

• sp3d2 hybridization

  

• octahedral

• sqaure pyramidal (1 lone pair)

• square planar (2 lone pair)

• t shape (3 lone pair)

• linear (4 lone pair)

explain

molecular polarity

• Polar Molecule: has a net dipole moment

• Nonpolar Molecule: no net dipole moment

• Compare bond polarity vectors in 3D

define

intermolecular forces

• Forces between adjacent molecules

• NOT forces within individual molecules

• Large impact on boiling points and melting points

types of IMFs

dispersion (van der waals)

dipole-dipole (van der waals)

hydrogen bonding

ionic interactions

explain

IMF: dispersion force

• Weak interactions between molecules that occur because of temporary distortions in the electron density of a molecule.

• All molecules experience dispersion forces

• Greater surface area = greater dispersion forces

• all compounds experience this to some degree

• temporary dipole (weak interaction between these dipoles is the force itself)

• 
  

• 
  

  

explain

IMF: dipole-dipole

• Attractive forces between two permanent dipoles of two polar molecules

• interactions happen when dipoles of adjacent molecules are aligned in a way such that the partially positive and the partially negative charges are in close proximity

• force is stronger than dispersion forces

explain

IMF: hydrogen bonding

• A strong attractive force between especially polar molecules…

• An H atom covalently bonded to a highly electronegative atom (N, O, or F) in one molecule interacts with a lone pair of electrons on N, O, or F in another molecule

• stronger than dispersion and dipole-dipole

explain

IMFs: ionic interactions

• electrostatic force between ions

• very very strong (stronger than H-bonding, dipole dipole and dispersion)

• ionic interactions result in bond strengths that exceed the strength of covalent bonds

effect of IMFs on boiling point

+melting point

• Temperature for liquid → gas phase change

• When substance’s vapour pressure equals the atmospheric pressure

• Boiling breaks intermolecular interactions

• Stronger intermolecular forces = greater energy required to overcome these forces = higher boiling point temperature

• Consider type of intermolecular force and surface area

melting point has the same trends as boilding points

• Temperature for solid → liquid phase change

explain

valence bond theory

• Explains discrepancies between VSEPR shapes and geometries of pure atomic orbitals

Three principles/assumptions:

1. Covalent bonds are formed by orbital overlap

2. Electrons are localized

3. Orbitals may hybridize before bond formation

principles of VB theory

1. covalent bonds between two nuclei are formed when each atom contributes one valence electron to a common orbital. this common orb., which is an overlap of atomic orbital, contains two electrons of opposite spins (sigma bond)

2. as a result of the overlap, the electrons are deemed to be localized. in other words, they are restricted to the area between two nuclei

3. elements of the second or higher-numbered rows of the PI can form bonds using pure atomic orbitals or “hybrid” atomic orbitals

4. hybrid orbitals, if they form bonds, always form sigma bonds

define

sigma bond

• a bond formed by the overlap of orbitals in a region of space that les on the same axis as the two nuclei (head on overlap)

• cylindrically symmetrical about the internuclear axis

• A single bond is a sigma bond

• Overlap of orbitals on axis with the two nuclei

• Atomic orbital – atomic orbital or

• Atomic orbital – hybrid orbital or

• Hybrid orbital – hybrid orbital
  
  

define

hybrid orbitals

• Mixing atomic orbitals to produce hybrid orbitals

• explains molecular geometry

explain

hybrid orb. sp3

• 4 hybrid equal orbitals (s and 3 p) = 3 degenrate hybridized orbitals

• four sp3 orbitals

• 109.5 degrees apart

• 4 regions of electron density

tetrahedral 109.5 degrees

trigonal pyramidal ( 1 lone pair) 107 degrees

bent (2 lone pairs) 104.5 degrees

explain

hybrid orb. sp2

• 3 pure atomic orbitals in (with one extra) = 3 equal degenerate hybridized orbitals …pure p left

• Three sp2 orbitals oriented

• 120° apart

• Three regions of electron density

trigonal planar 120 degrees

bent >120 degrees

explain

hybrid orb. sp

• 2 pure orbitals in (with two p left) = 2 equal degenerate hybrid orbitals

• Two sp orbitals oriented 180° apart

• Two regions of electron density

linear - 180 degrees

explain

hybrid orb. sp3d

• 5 sp3d orbitals oriented 120°/ 90° apart

• Five regions of electron density

trigonal bipyramidal

seesaw (1 lone pair)

T shaped (2 lone pair)

linear 3 lone pair

explain

hybrid orb. sp3d2

• 6 sp3d 2 orbitals oriented 90° apart

• six regions of electron density

octahedral

square pyramidal (1 lone pair)

square planar (2 lone pair)

T shape (3 lone pair)

linear (4 lone pair)

explain

double bonds ( in reference to hybridization)

• Consider ethene (ethylene)

• Each carbon atom has three regions of electron density and is sp2 hybridized

• Sigma (σ) bonds form from overlap of orbitals on axis (head-on overlap)

• But each C has a 2p orbital remaining!

• Sideways overlap of 2p orbitals gives a pi bond

• Recall that each 2p orbital has two “lobes”

• Due to sideways overlap, pi bond cannot rotate

• 1 double bond = 1 σ bond + 1 pi bond

explain

isomers due to a double bond

• σ bonds can rotate, but pi bonds cannot

• cis/trans isomerism is a result of pi bonds not being able to rotate

• cis/trans isomers have the same formula, the same atom connectivity, but a different spatial arrangements of the groups at the C=C

explain

triple bonds ( in reference to hybridization)

• Consider acetylene:

• Each carbon has two regions of electron density and is sp

• sigma bonds formed by direct, on-axis overlap of orbitals

  

  
  

• Each carbon has two 2p orbitals remaining... they are perpendicular to the sp hybrids

• 1 triple bond = 1 σ bond + 2 pi bonds