CSU’s Gen Chem 111 Test 3 – Flashcards
Unlock all answers in this set
Unlock answers| what forms ionic bonds? |
oppositely charged ions |
| what forms covalent bonds? |
| positively-charged atomic nuclei and the negatively-charged electrons between them |
| what does chemical bonding do? |
| Chemical bonding lowers the potential energy between positive and negative particles |
| nonpolar covalent |
| no electronegativity difference |
| polar covalent bond |
| less than 1 |
| ionic bond |
above 1 electronegative difference metals and nonmetals |
| types of covalent bonds |
nonmetal and nonmetal metalloid and nonmetal |
| what is bond polarity? |
a measure of how equallyelectrons are shared |
| what happens in a completely nonpolar covalent bond |
the electrons are shared equally ;;EN = 0 |
| what happens in a nonpolar covalent bond? |
;The electrons are shared fairly equally between the two atoms that are bonded, and ;EN ; 0.4 |
| what happens in a polar covalent bond? |
;The electrons are pulled more strongly to one atom than to the other atom, and ;EN = 0.4 ; 1.7 |
| what happens in an ionic bond? |
;The electrons are transferred ;;EN ; 1.7 |
Which is the most polar bond?
(a) N - H ? (b) F - N ? (c) O - H ? (d) I - Cl |
| O-H |
| what makes up a lewis electron dot symbol? |
•consists of the chemical symbol for an element surrounded by dots • Chemical symbol represents nucleus + core electrons • Each dot represents a valence electron
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| why do we use lewis dot diagrams? |
| used to show and track the valence electrons (e.g. during a chemical reaction) |
| how to build a lewis diagram |
•Figure out the total # of valence electrons.
• Make sure to adjust if it’s a cation (remove electrons) or an anion (add electrons). • • Figure out how the atoms are connected. • The central atom is generally less electronegative than the atoms around it (except for H). • Start by connecting them all with single bonds. • • Complete the octets of the atoms bonded to the central atom first (but for Hydrogen only use 2 electrons). • • If there aren’t enough electrons for the central atom, try giving it multiple bonds. • • Place any left over electrons on the central atom.
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| What is the octet rule? |
•Octet = eight valence electrons = four pairs of valence electrons = eight v. e. means full s and p subshells = filled outer level
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| what is formal charge? |
| It’s the charge that an atom would have if all the bonding electron pairs were shared equally. |
| what is the formal charge equation? |
| Formal charge = (# of valence electrons in isolated atom) – (# of electrons assigned in the Lewis structure) |
| what is the most stable lewis structure? |
| 1) has atoms that bear formal charges closest to zero, and (2) any negative formal charges are on atoms that are more electronegative. |
| what is more important: the octet rule or formal charge |
| satisfying the octet rule is more important than minimizing formal charges. |
| exceptions to the octet rule |
•Odd number of valence electrons in molecule •Less than an octet (H, Be, B, Al common) More than an octet (n = 3 and above since d orbitals available |
| what is a resonance structure? |
describing Bonds in the molecule have: |
| what does bond order equal? |
atom sets |
| what does VSEPR stand for ? |
| Valence-Shell Electron-Pair Repulsion Theory |
| what is vsepr? |
| Model explaining/predicting the geometric shape of molecules: Electron ‘groups’ are as far apart as possible Requires an appropriate Lewis structure Number of bonding groups and nonbonding groups are arranged in optimal geometric arrangements |
| what determines the electron group arrangement? |
| total number of groups. |
| what are the types of groups in molecules? |
| Bonding groups (X) = # of surrounding atoms (regardless of bond order) • Nonbonding groups (E) = # of lone pairs |
| two electron groups: |
electron group arrangement: linear Molecular shape (AX2): Linear |
| three electron groups |
electron group arrangement: trigonal planar Molecular shape (AX3): Trigonal planar Molecular shape (AX2E): Bent (V shaped) |
four electron groups
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electron group arrangement: tetrahedral Molecular shape (AX4): Tetrahedral Molecular shape (AX E): Trigonal pyramidal 3 bonded atoms Molecular shape (AX2E2): Bent 2 bonded atoms |
| five electron groups |
electron group arrangement: trigonal bipyramidal Molecular shape (AX5): Trigonal bipyramidal |
| Five Electron Groups Electron-group arrangement: Trigonal bipyramidal |
Molecular shape (AX4E): Seesaw Molecular shape (AX E ): T-shaped Molecular shape (AX2E3): Linear |
| Six Electron Groups Electron-group arrangement: Octahedral |
Molecular shape (AX6): Octahedral Molecular shape (AX E): Square pyramidal 5 bonded atoms Molecular shape (AX4E2): Square planar 4 bonded atoms |
| summary of vsepr theory (steps) used to predict chemical properties |
molecular formula lewis structure (count all e groups around central atom) electron-group arrangement (note positions of any lone pairs and double bonds bond angles (count bonding and nonbonding e groups seperately) molecular shape |
| valence bond theory |
uses wave behavior of the electrons to explain bonding Bonds form when the orbitals of two atoms that why H-H forms H2 |
| Bond strength and orbital overlap |
| greater the orbital overlap - the stronger the bond |
| overlap depends on : |
| shape and direction of the orbitals |
| bonds are oriented: |
| in the direction that maximizes the overlap |
| hybridized atomic orbitals |
| Mathematically ‘mix’ isolated atomic orbitals to obtain hybrid orbitals |
| number of atomic orbitals = |
| number of hybrid orbitals |
| type of hybrid orbitals depends on types of |
| atomic orbitals mixed |
| Sigma (σ) bonds: |
| result from end-to-end overlap of orbitals – Produces a region of high electron density directly along bond axis |
| Pi (π) bonds: |
| result from side-to-side overlap of orbitals – Produces two regions of electron density above and below bond axis |
| bond order and types: |
| Single bond: one σ bond • Double bond: one σ and one π bond • Triple bond: one σ and two π bonds |
| Determining Geometries of Axn Molecules |
| 1. Draw the Lewis Structure 2. Determine the total number of electron domains around the central atom A. 3. Determine the electron domain geometry (arrange the electron domains so that repulsions among them are minimized). 4. Use the resulting arrangement of the bonded atoms to determine the molecular geometry. |
| How the Type of Electron Pair Affects Bond Angles |
| Repulsive force of electron domains/volume occupied by electron domains: Nonbonding > triple > double > single pairs bonds bonds bonds |
| Nonbonding pairs experience less blank blank than bonding pairs |
| nuclear attraction |
| multiple bonds contain higher blank blank blank compared with single bonds |
| electronic-charge density |
| memorize table 11.1 |
| physical properties of sigma and pi bonding |
Sigma bonds allow free rotation of the atoms around the bond axis
Pi bonds restrict rotation around the bond axis (π orbitals must |
| macroscopic observations of gas |
conforms to shape and volume of container high compressiblilty high ability to flow |
| macro observations of liquid |
conforms to shape of container; volume limited by surface very low compressibility moderate ability to flow |
| macro observationsof solid |
maintains its own shape and volume almost none compressibility and ability to flow |
| Intramolecular forces (bonding forces) |
| These forces exist within each molecule. – They influence the chemical properties of the substance. |
| Intermolecular forces (nonbonding forces) |
| These forces exist between molecules. – They influence the physical properties of the substance. |
| phase changes require what |
| changes in the energy of matter (You are not breaking bonds – just the weaker intermolecular forces break) |
| sublimation |
| solid to gas |
| deposition |
| gas to solid |
| comparison of bonding forces |
Ionic: cation to anion covalent: nuclei shared e- pair metallic:cations delocalized electrons |
| comparison of nonbonding (intermolecular) forces |
ion-dipole: ion charge-dipole charge h bond: polar bond to H-dipole charge (high EN of N,O, and F) dipole-dipole: dipole charges ion induced dipole: ion charge - polarizable e cloud dipole - induced dipole: dipole charge - polarizable e cloud dispersion (london): polarizable e clouds |
| bonding and nonbonding forces are just electrostatic attractions. therefore: |
| Bonding forces tend to be stronger (large charges - close together) – Nonbonding forces tend to be weaker (partial charges – further apart) |
| ion- dipole forces |
| Most commonly occur in ionic solutions (i.e. salt water) – Charge of the ion and partial charges of the polar molecules are attracted – In salt water – ions are solvated by water (hydration) |
| dipole-dipole interactions |
| Dipoles in polar molecules are attracted • Attraction creates directional orientations in these substances (oppositely-charged poles point at each other) |
| why is hydrogen bonding a special case? |
H-bonding is a dipole-dipole force where there is an attraction between a H H-bonding is only present in molecules |
| charge induced dipole forces |
| Polarizability - distortion (or squishiness) of an electron cloud • Increases down a group - size increases and the larger electron clouds are further from the nucleus • Decreases left to right across a period - increasing Zeff shrinks atomic size and holds the electrons more tightly |
| dispersion (london) forces |
Instantaneous dipoles caused by the random motion of Stronger forces exist between molecules with more ‘surface area’ |
| ions present in these particle forces |
ionic bonding (ions only) ion dipole forces (ion + polar molecule)
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| ions not present in these interacting particle forces: |
dipole dipole forces (polar molecules) polar and nonpolar: dipole induced dipole forces nonpolar molecules only: dispersion forces only |
