Structure 2.1-2.2

• Ionic compounds: nearly always solids ("salts" - if liquid then melting point is under 100 Celsius) because of their high melting points, low volatility, high conductivity when molten of dissolved in solutions, with high brittleness with strong bonds that fracture easily with the slightest shift in layers

  
  

  Made of metals + semi metals (ex. NaCl)
• Covalent compounds: usually gases or liquids at room temperature (if solid then melting point above 350 degrees celsius), volatile, nonconductors, weak intermolecular forces between molecules so they are soft and malleable
• Molecular compounds: bonding of two non/semi-metals to form a compound
• Chemical properties are affecter by electrons
• Physical properties are affected by neutrons
• Equivalence point is the point at which the moles of both the acid and base are in equal amounts (neutralization), usually measured visually through the use of an indicator.
• Moles (mol) are groups of particles = 1mol =6.02x10^23 particles
• Molecular mass is the combined atomic mass for the molecules in a compound (ex. H2O = 2(1.01)+16=18.02g/mol)
• Formula units are the units of the formula of a compound. They are the same as the number of atoms of each element in the compound for ions
• Allotropes are different forms of an element in the same physical state
• Alloys are mixtures created through metallic bonding
• Isotopes are different variations of a singular, identical element with different masses
• Elements are building blocks of matter with only one type of atom, unable to be broken down further
• Random error is an error in precision that is accounted to flaws in tools that make results change completely
• Systematic error is an error of accuracy where there are flaws in the methodology, causing the obtained values to be close to the predicted but over a wide range
• How can we find the group of an element?

  On its emission spectrum, count the number of electrons lost before a huge rise in ionization energy which signifies all valance electrons are lost
• Transition metals tend to lose their 4s electrons first when it comes to ionization
• How do we identify the moment at which an atom has been ionized?

  On the emission spectrum, it is the first big gap in emission that signifies a large rise in ionization energy
• The normal order for electronic configuration goes as follows: 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s- increases by PERIODS where s orbitals can have a power up to 2, p orbital up to 6, and d orbitals up to 10
• 2 exceptions to normal electronic configuration?

  Chromium (Cr): [Ar]4s^1 3d^5

  Copper (Cu): [Ar] 4s^1 3d^10
• The energy level from any level away from the nucleus emitting back to n=3 causes Infrared waves to be emitted.

  
• The energy level from any level away from the nucleus emitting back to n=1 (ground level) causes ultra-violent waves to be emitted.
• The energy level from any level away from the nucleus emitting back to n=2 causes visible light waves to be emitted.
• Intermolecular bonding is only for non-metals- they keep compounds together by attracting the molecules to each other- like dipole dipole forces, London Dispersion forces, and Hydrogen bonding. Intramolecular forces are used to keep molecules together- for metals, it can be metallic bonding, others it can be ionic or covalent bonding.
• Hydrogen bonding: strongest intermolecular force, hydrogen is strongly attracted electronegative elements like oxygen, fluorine, and nitrogen and bonds to their lone pairs
• Dipole-dipole forces: 2nd strongest intermolecular force that is stronger in more polar molecules
• London Dispersion forces: weakest intermolecular force, non-polar force but present in ALL molecules, stronger the heavier the molecule is. Usually in gas form
• Dipole-dipole induced force: An intramolecular force using delta signs for positive and negative dipoles (terminals)- it is the creation of temporary dipoles due to electrons' rapid, random movements that might cause a slight higher concentration at one end, giving the atom negative and positive dipoles.

  
• Resonance structures displace a double bond in an attempt to instantly quell any electronic irregularity (double bond moving to an electron deficient carbon)
• The main resonance contributor (MAJOR) is the resonance structure which is the most stable
• Tertiary carbo-cations are more stable than secondary carbo-cations which are more stable than primary carbo-cations

  
• As carbon-cation degree increases, more methyl groups are added- electron-donating groups. These groups give electrons to the central carbon through either hyper-conjugation or inductive effect which balances the cation's positive charge and causes it to be more stable
• Hyperconjugation

  The pass of electrons through the overlapping of orbitals
• Inductive effect

  The passing of electrons through sigma bonds
• When picking the Major resonance contributor consider

  1. Atomic size (bigger is more stable to shift charges to- diluted effect)
  2. Valance shell + shared electrons (structure that fulfills octet rule perfectly will be more stable)
  3. Electronegativity (more electronegative elements are more stable with negative charges/more electrons)
  4. Degree of carbon-cation (3>2>1)
  5. Priority in resonance (which of these factors are of priority to the molecule?)
• Metalloids show intermediate properties between metals and non-metals
• Group 4 have no specified oxidization state as they have a low tendency to form ions
• Transition metals always tend to lose from the 4s sub-level first before 3d
• Zn is not a transition metal as it cannot form colored complexes due to its lack of multiple oxidation states and full d-orbitals- its only oxidation state is +2
• Polyatomic ions are usually formed through 1 ionic bond and 1 covalent bond