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Electromagnetic radiation is the emission and abosrption of energy in the form of electromagnetic waves Increasing wavelength goes left to right, increasing light energy goes right to left - inverse relationship Wavelength is the distance between 2 peaks or troiughs (bottoms) of the waves Frequency(hz) is the number of wavelengths per second that travel past a certain point All Electromagnetic radiation travels at the speed of light Product of frequency and wavelength always equals a constant, c, the speed of light Could be in m/s or nm/s two different measurements As frequency increases, wavelength decreases, inverse relationship Inverse: as one increases, the other decreases Direct: as one increases, the other increases (use simple math to prove it) Quantum: smallest quantity of energy that can be emitted or avbsorbed in the form of electromagnetic radiation Energy is quantized: only certain energies are allowed Energy must have a certain value to exist For electrons to be in a certain energy level they must have certain energy Plancks equation: E = hv E is energy in joules H is plancks constant (6.626e-34 js or kgm^2/s) v = frequency Energy can only be gained or lost in whole number multiples Electromagnetic radiation itself is quantied and can be viewed as a stream of particles called PHOTONS Photon: a quantum of electromagnetic radiation Energy is inversly related to wavelength but directly to frequency
All waves have the same velocity because it represents electromagnetic radiation and all electromagnetic radiation travels at the speed of light Wavelength is inverse to energy so smaller wavelength is bigger energy Einstein and planck proved electromagnetic radiation exhibits wavelike and particulate properties (wave particle duality) Debroglie then proved matter exhibits particulate wavelike and particulate properties Continuous spectrum White light passed thru a prism produces a spectrum of color The electron in hydrogen was circling the nucleus of the atom in circular orbits that are quantized The further you go, the higher the energy of the orbit the higher the energy of the electron The energy of the electron in Energy levels: n=1, n=2, … n=1 is the ground state - stable N = quantum number n >1 is unstable n=infinity - electron leaves atom The wavelengths of light are occurring because of electron energy levels The lines of colors are caused by electron dropping from outer orbits to the n=2 orbit An electron in n=1 can get excited from energy that we calculated in E=hv When it releases the correct amount of energy it will move back down to n= Delta E is the change in energy When energy is emitted DE is negative, when gained its positive The energy change when the electrom in a hydrogen atom moves from one energy level to another can be calculated with (check slide for formula) This equation can only be used on one electron species(like hydrogen), not multi electron species E = positive value of DE Line spectra: each element generates unique lines of color or emission spectrum In bohrs failure
M provides maximum number of orbitals for different values of l We dont need to know which p orbitals have which ml values Just types of ml values The spin quantum number can only have values of +- 0. Orbital will become unstable if e- have the same spin No 2 e- in an atom can have the same 4 quantum numbers 2 e- in an orbital have different ms values 2 e- per each orbital w opposed spins For 2 e- to be in an orbital they must spin in opposite directions (pictured by the arrows) E- do not pair in an orbital until all other orbitals in the subshell have been occupied by a single e- (you cant be married before being single) electron configuration exceptions: Chromium: not predicted by hunds rule It is more stable if 1 e- is in an orbital vs if there are 2 e- in another orbital but it is less stable Copper: not predicted by hunds rule When the 3d shell isnt full the atom is less stable then if it is full Memorize these 2 differences Excited states To write the electron configuration of an excited state you must excite the n value E- from the s orbital are removed before the d orbital Only used for transition metals
Elements in the same groups have similar chemical properties Inner shell e- are core e- and interfere w attraction between nucleus and valance electrons Using bohrs model to draw e- (not correct just for drawing purposes) 2 inner electrons: shield the 2s electrons from nucleus Repel each other as well as the 2s electrons Also being pulled in by the Zeff Pulling and pushing happening at the same time to keep them in place Bc of the shielding the 2s e- are not feeling the full +3 charge That diminished charge is called Zeff and is how cations and anions are formed bc the electrons can be pulled away easier Every inside shell acts as a shield for the nucleus, meaning the outer shell is weaker and e- can be pulled away 1s e- are core electrons and feel the full full pull and are super hard to pull away They have full attraction, no shielding Closer to the nucleus is greater attraction, farther is less attraction I1 is the energy needed to remove the highest energy electron (electron farthest away from nucleus) of a gaseous atom and move it farther away The value of Z effects the ionization energy z=atomic number So does the value of Zeff The core electrons help for 2s electrons to be removed easier bc they are shielding electrons Periodic trend anomalies: From Be to B there should be an increase but theres a decrease Same for N to O, Mg to Al and P to S its because they are right on top of each other Full e- shells are super stable (2 e in 1 square) For Be to B: It takes less energy to pull away an e- in a half full shell than a full shell which is why theres a drop between certain elements, because its easier to remove unfull subshells For N to O: when a second electron is added to the p orbital in oxygen, the e- repelling each other makes it easier to remove an e-, causing the drop in the trend WILL BE TESTED ON THIS
4 types of bonds Intramolecular forces = chemical bonds, hold groups of atoms or ions together and make them function as a unit Form molecules, compounds, etc Ionic Covalent bonds: Polar covalent bond Non-polar covalent bond Metallic bonding Dont need lewis structures for transition metals Ionic bonding: metal and nonmetal bond Metal looses e- to form a cation Nonmetal gains an e- to form an anion E- transfer takes place (e- goes from metal to nonmetal) Memorize list of metals that form multiple cations A lot come from transition metals Especially: silver, zinc and cadmium - they dont form multiple cations, they only form these special ions If you have multiple ions you use roman numerals Hg is diatomic and always occurs together to form ions Know all polyatomic ions on list in notes Covalent bonds have only nonmetals No electron transfer, electrons are being shared Not talking about ions at all Difference between ionic and covalent is that ionic give each other electrons to bond vs covalent the e- are shared Non-polar covalent bonds: electrons are equally shared Lone pairs: e- that arent being shared Bonding pairs: e- being shared Why do atoms want to bond? As atoms come together, energy drops to remain stable The energy that forms when bonding occurs is why atoms bond Perfect distance - atoms are bonded but not too close - referred to as bond length When getting too close, energy increases because of too much repulsion
We want low energy for atoms when bonding We form bonds because it leads to lowering of energy Why do we want to lower energy? Memorieze prefixes for namiing covalent compounds to tell ppl how many atoms are in the compound When element begins with a vowel take off ending o or a Peroxide is 02 2- Memorize the peroxides In ionic compounds, dont include things like di in dioxide Polar covalent bonding: e- are not shared equally Because of that we develop charges on each species (like a partial negative or a partial positive charge) Elements that can pull e- better have a partial negative charge Electronegativity: periodic trend, ability of an atom to attract e- in a chemical bond (ability to be more or less polar) Get table on constant sheet showing the electronegativty of each element Greater EN distance between atoms, the greater the polarity Check notes for breakdown Metallic bonding: metal and metal bonding Delocalized - because e- are farther from nucleus e- can move around
Shapes of simple molecules (shapes with central atoms and terminal atoms, cannot talk about shapes for complex atoms) Lewis structures in different orientations Dipole moments: the separation of the delta charge (seperation between delta positive and delta negative) Present in polar molecules Symmetrical and asymmetrical molecules Polar molecules are asymmetric Linear: set in stone shape, 2 pairs of e- (double or single bonds) on the central atom, 2 atoms surrounding it, 180 bond angle If no lone pairs on central atom, molecule shape = molecule geometry Vspr: valance shell e- pair repulsion To figure out shape: treat double bonds as single bonds Triganol planar: 3 e- pairs on central atom, no lone pairs, 120 bond angle Bent: 3 e- pairs, 1 is a lone pair, bends other atoms because e- are repelling each other, lone pairs repel bonding pairs, has same bond angle as triganel planar Molecular geometry is the shape it holds, molecular shape specifies lone pairs or not Bent is always polar (?) tetrahedral : 4 e- pairs on central atom, 109.5 bond angle Triangonal planar: 1 lone pair 3 e- pairs, 109.5 bond angle aprox. Because lone pairs are pushing down on the bonds, moving them, polar Lone pairs push e- away or repel them, and in doing so push e- into the orientation of the shape with the same amount of e- pairs including lone pairs, but that is the geometry, with lone pairs it assumes a new shape Bent: 2 lone pairs, 2 e- pairs, < 109.5 aprox angle, polar Trigonal bypyramidal: 5 e- groups, 180, 90 and 120 bond angles (because there are 2 pyramids)
Seesaw: 1 lone pair, 4 e- groups, 180 bond angle (90 and 120 gone bc lone pair is not a real atom and therefore there is no angle between that and other molecules), polar T shaped: 2 lone pairs, 3 e- pairs, 180 bond angle, polar Linear: 3 lone pairs, 2 e- pairs, non-polar, 180 bond angle Octahedral: 6 e- pairs, 180 and 90 bond angles Square pyramidal: 1 lone pair, 5 e- pairs, polar Square planar: 2 lone pairs, 4 e- pairs, non-polar, 90 bond angle
Carbon combines its 2s and 2p orbitals to form four new orbitals called sp3 hybrid orbitals Sigma bonds are single bonds To get 2sp we combine only 2 of the p orbitals and the s orbital One p orbital is left over and will be left as unhybridized Helps form a double bond or a pi bond Can only be formed on overlapped hypbridized orbitals Unhybridized p orbitals are parallel and if they arent, pi bonds will break Single bond is formed by overlapping hybridized orbitalsm pi bonds formed by unhybridized orbitals For linear shapes, use the 2s orbital and only 1 2p orbital, 2 sp orbitals formed, 2 p orbitals left unhybridized For a triple bond you have 2 pi bonds bc of 2 unhybridized p orbitals, and 1 sigma for 1 hybridized p orbital Must have the same orientation for maximum overlap so bonds wont break Howcome we can do resonance structures which switches orientation of molecules when pi and sigma bonds prevent it? Cannot put lone pairs in a double bond Can bring 1 d orbital e- to mix with sp orbitals Unhybridized are higher in energy than hybridized
Intermolecular forces London Dispersion forces Dipole dipole interactions Etc London dispersion forces All atoms have them bc all atoms have electrons, the forces occur when electrons gather Gathering of e- is instantaneous dipole moments, happens and is gone instantly Weak and short lived forces between molecules Ca has more e- than ca2+ amd it is longer than ca2+ so e- can be easily distorted as they are not under the + pull of rhe nucleus as much on ca2+ Small species means you are less mobile, need to be mobile for polarizability A pi bond helps e- move around Sigma bonds cause e- to be stuck between multiple atoms Dipole- dipole interactions must be polar Higher boiling point needs greater attraction between molecules List intermolecular forces, whoever has more wins ion-dipole interactions An ion and a polar molecule Water (polar) separates and ionic molecule by weaving into its lattice structure The delta negative part of water (O) is attracted to the positive ion and vice versa Whichever has the highest ion has the highest ion-dipole interactions Hydrogen bonding Strong dipole-dipole force between the H atom in a polar bond (F, O or N) and a non-bonding electron pair (Lone pair) o a nearby electronagitave ion or atom Interaction between the lone pair on F, N, or O (super electronegative) and Hydrogen dipole-dipole force Delta - to delta + matching between molecules Identify the correct H (connected to F, N, or O), then find F, N, or O with a lone pair, then bond Hydrogen bonds need to be kinda close Surface tension, adhesion, capillary action, and cohesion all relate to hydrogen Vapor pressure has a inverse relationship to imfs
Do not rush Limiting reactant must be completely used by the end of rxn Determines amount of product formed Determines how much of other reactasnts are used
Something that is polar will dissolve something that is polar
Lattice energy of an ionic compound: on final (worksheet 11 pt 2) Bond energies will always be positive bc eg is needed to go into bond to break it Must be in gas phase Forming bonds releases energy Breaking bonds needs energy Must break bonds to form bonds in balanced equations Break reactant bonds and form product bonds We count a bond as a mol of that type of bond CH4 has 4 mols of C-H bonds Double bonds count as 1 bond but are specified that they are double bonds O- - O bonds Combustion rxns are exothermic and have -4H E- transfer is endothermic and needs eg so it isnt favorable Ionic compounds are formed because the ions created during the electron transfer when they come together they release more energy than the e- transfer and is very favorable Lattice energy problems are similar to hess law problems - flipping, multiplying and canceling out equations Use lattice energy periodic trends formula to answer theoretical questions, not mathematically solve it
