Chem Study Guide PDF

- Basic Calculations: simpler math we’ve done, not necessarily related to specific content ▪ Dimensional analysis ▪ Calculating Molarity, M ▪ Calculations involving dilutions of solutions, M1V1 = M2V2 - Absorbance Spectroscopy: See the two LabPa...

- Basic Calculations: simpler math we’ve done, not necessarily related to specific content ▪ Dimensional analysis ▪ Calculating Molarity, M ▪ Calculations involving dilutions of solutions, M1V1 = M2V2 - Absorbance Spectroscopy: See the two LabPal Prep Activities and Lab Activities we did ▪ Beer’s Law A=bc o What does tell us about the absorbing molecule? - This is a measure of how much light is being absorbed by the solution o Plots of Beer’s Law data ▪ What data goes on the “y” axis? - Absorbance ▪ What data goes on the “x” axis? - Concentration ▪ How can you use the fit equation y = mx + b? What data can you determine using that equation, and what data would you use or “plug into” that equation? Y= absorbance/ slope= molarity absorbance/ x= concentration Chapter 2 Basic Atomic Structure ▪ Nucleus has protons + neutrons, this sum gives you atomic mass ▪ The rest is electrons, makes up most of volume ▪ # of protons is what determines an element Chapter 3 ✓ Wave Nature of Light ▪ Wavelength, frequency, speed of light, amplitude, nodes ▪ Know relative ordering of visible light and wavelength range o long wavelength to short wavelength: ROYGBIV - E = lesser, then the wavelength is longer  o Know the wavelength range, ~700 to ~400 nm in order from red to violet - Red: 630 nm - 700 nm - Orange: 590nm - 630 nm - Yellow : 550 nm - 630 nm - Green: 490 nm - 550 nm - Blue: 440 nm-490 nm - Purple: 400- 440 nm o I will provide a visible spectra like I did on Quiz if needed ✓ Particle Nature of Light ◼ Planck & Einstein: light is quantized: o Energy of a photon = h= hc/λ remember this has units of J/photon ◼ Atomic Emission Spectra – o What it is? What “causes” it? Why is it different for different elements? - What is it:An atomic emission spectrum is a spectrum of the electromagnetic radiation emitted by an atom or ion when it transitions from a higher energy state to a lower energy state. The spectrum consists of distinct lines, each corresponding to a specific wavelength of light emitted by the element, representing the energies of the photons released. - What causes it: - Electron excitation: when atoms absorbed energy - Energy release: As these excited electrons return to their ground state, they release energy in the form of light. - Quantized Energy levels: The energy levels of electrons in an atom are quantized, meaning they can only occupy specific energy states. This quantization leads to the emission of specific wavelengths of light, resulting in a line spectrum. o Bohr’s Model of H atom - Bohr model: electrons in atoms have fixed energies relative to nucleus o How to calculate Energy, wavelength, frequency or n values associated with transitions in H atom – I will give both equations associated with transitions in one electron species o See the Electronic Transitions in Hydrogen Atom LabPal - Bracket series E(n) to E(n=4) - Visible nf = 2 ✓ Electron configurations – ▪ Pauli’s Exclusion Principle, Hund’s Rule, Aufbau Principle- - Pauli's Exclusion Principle: - The Pauli Exclusion Principle ensures that each electron in an atom has a unique quantum state. - Hund's Rule dictates how electrons are distributed in degenerate orbitals, promoting stability through maximum spin multiplicity. - The Aufbau Principle outlines the order in which electrons fill atomic orbitals, starting from the lowest energy levels. o Be able to use/apply each ▪ Be able to write an electron configuration in either of two ways we looked at for ANY element o Complete notation – That is, what element is 1s22s22p63s23p3? o Abbreviated notation – The above is equivalent to [Ne]3s23p3 ▪ Be able to understand what an Orbital Box Diagram is showing ✓ Be able to connect a Photoelectron Spectrum for an element to the electron configuration ▪ Understand how electron configurations lead to chemical properties - Electron configuration: The distribution of electrons in an atom's orbitals determines its electron configuration ▪ Understand how & why elements with similar electron configurations (elements in same groups) have similar chemical behavior  - Isoelectronic- having the same electron configuration - Main group metals: lose all their valence electrons - Losing electron from highest “n” level - If toy have more than one “L” sublevels within an “n” level, lose from highest “L” first - Transition metals: lose from highest “n” level 1st - Nonmetals: gain enough electrons to be like noble gas following them on PT - Ionic bonding: metals that lose electrons to become cations react with nonmetals that gain those electrons to become anions ✓ Periodic trends –These make good multiple choice questions ▪ Zeff can help us understand these – so what is that? What is trend in Zeff ? - Zeff: the net positive charge experienced by an electron in a multi-electron atom, accounting for the shielding effect of inner electrons. ▪ Atom size even easier to understand and use to remember other trends – so what are the trends in the size of atoms? ▪ In atomic size/atomic radii - Increases down and to the left - Atomic size gets smaller as Zeff increases - Metals tend to be longer - Since metals are longer, they will lose electrons easily - Nonmetals tend to be smaller - Since they are smaller they gain electrons more easily ▪ In ionization energy–How easy or difficult it is to remove an electron - High IE means it is difficult to remove electron - Trend: increases up and to the right ▪ In electron affinity – How willing an atom is to take on more electrons - High EA is a more negative EA, atoms with high EA want more electrons ▪ In ion sizes - Cations are smaller than their parent atoms, anions larger than parent atoms - Cations have less electrons with the same number of protons - Anions have more electrons than protons ▪ In electronegativity– How strong an atom pulls on electrons in a bond - High EN means atom is pulling strongly on electrons in bond Parts of Chapter 4 Chemical Bonding and Chapter 21 Organic Chemistry ✓ Ionic Bonding - The transfer of electrons from one atom to another ▪ How are electrons involved in this type of bonding? (they’re transferred) - cations : metals / few electrons in their outer shell, tent to lose one or more electrons to achieve a stable electron configuration  - Anions formations: nonmetals: more electrons in their outer shell and are closer to achieving a full valence shell, tend to gain one or more electrons ▪ What types of atoms bond this way? - Metals and nonmetals ▪ How is this type of bonding related to periodic properties? (metals lose electrons because they have low IE, low EA, so what do nonmetals do and why?) - Metals: - Low Ionization Energy (IE): Metals have low ionization energies, which means they can lose electrons relatively easily. This property allows metals to form cations and participate in ionic bonding. - Low Electron Affinity (EA): Metals have lower electron affinities compared to nonmetals, meaning they are less inclined to gain electrons. This reinforces their tendency to lose electrons. - Nonmetals: - High Ionization Energy (IE): Nonmetals have high ionization energies, making it more difficult for them to lose electrons. As a result, they are not likely to form cations. - High Electron Affinity (EA): Nonmetals have high electron affinities, which means they are very likely to gain electrons and become anions. This characteristic is crucial for their role in ionic bonding. ✓ Naming Ionic Compounds – See the Handouts ▪ Binary Compounds – just two ions ▪ Compounds that contain Polyatomic Ions – you do need to memorize these ▪ Acid Names from Anion Names ✓ Covalent Bonding – ▪ How are electrons involved in this type of bonding? - (they’re shared between atoms) ▪ What types of atoms bond this way? - (nonmetals) ▪ How is this type of bonding related to periodic properties? - (nonmetals share electrons in covalent bonding because they have high IE, high EA) ✓ Covalent Bonding in the context of Organic Molecules – See the Molecular Structure Activities in LabPal o HONC 1234 – this mnemonic gives the preferred bonding patterns for H, O, N and C that we can expect to see in covalent molecules, mostly but not limited to organic molecules ▪ H always forms one bond ▪ O prefers to form two bonds, and will thus have two lone pairs ▪ N prefers to form three bonds, and will thus have one lone pair  ▪ C always always always forms four bonds! See below for more detail. o C always requires four bonds: can form 4 singles, 2 doubles, 2 singles and 1 double, 1 single and a Triple ▪ Given a molecular formula for a hydrocarbon (with or without O or N), be able to draw isomers using Lewis structures or Line diagrams - Eg C6H10O ✓ Isomers = same formula, different arrangement of atoms ▪ Given a condensed formula, recognize that this gives you the order to the bonding, Eg CH2CHCH2CH2COCH3, be able to draw Lewis Structure or Line diagram of the corresponding molecule ▪ Given a line diagram, be able to translate to molecular, condensed or structural formula (Lewis structure) ✓ Lewis Structures: Know, understand and be able to apply ALL BASIC RULES for drawing Lewis - Structures. In addition to the basic rules, - Carbon binds to 4 with no LP - Hydrogen bonds to 1 - Oxygen binds to 3