Electrons in Atoms So why does potassium explode in water? Quantum Mechanics Periodic Trends Chemical Bonding
12.1 Development of Atomic Models Dalton s Thompson s Rutherford s Bohr s carbon Quantum Model of helium shows 2 electrons in the 1s orbital or 1s 2 Bohr proposed electrons orbit in paths of fixed energy called energy levels.
Size of Atom and Subatomic particles
Quantum Theory Symphony
Quantum Mechanical Model The quantum model of the atom is based on the solution to the Schrödenger equation. One way to visualize the model s electron levels is to imagine a ladder where the higher rungs or levels are closer together. 7 rungs = 7 energy levels Principle quantum numbers (n) correspond to the energy levels
Atomic Sublevels or S and P orbitals Probability cloud models (left) show where it is most likely to find electrons. For each principal quantum number (n) there are the same number of sub levels. (n=2, 2 sublevels) Helium n=1, had 1 suborbital the s orbital. Level 2, n=2, has 2 sub levels, s and p. (The px,py and pz orbitals are found on energy levels 2-7) Each energy level, like level 2 shown here in green. has a spherical s orbital.
Ask Your Neighbor: 1. How many sublevels are there for each principle quantum number? 2. Which energy levels contain px, py and pz orbitals? 3. How are s orbitals different than p orbitals?
Electron Configuration of Hydrogen 4s 3 d 3s 3p How are all group IA elements similar in electron configuration? 2s 2px 2py 2pz 1s H 1s 1
Light will bend and reflect at the interfaces between different materials. Prism lenses bend light. White light is a blend of all wavelengths of visible light.
Each color has a different wavelength (λ) lambda = wavelength.7µm.6µm.5µm 1000 µm = 1 mm.4µm
The shortest wavelength also has the highest energy, hence UV light can harm us if the wavelength is too short! http://science.hq.nasa.gov/kids/imagers/ems/visible.html
Evidence of Energy Levels and Suborbitals http://phet.colorado.edu/new/simulations/sims.php?si m=neon_lights_and_other_discharge_lamps What are emission spectra? How is each element s emission spectra unique?
Why Discrete does each lines element = quanta have its own of energy signature emission spectrum? Tell neighbor. A. Each element has a different number of protons and electrons. B. Each element has unique nuclear attraction for electrons in shells. C. Each atom s first energy level is a unique distance from nucleus. D. Distance between outer energy levels in atom is unique to each element. E. Electrons emit photons whose frequency is proportional to energy lost.
Click here Predictions Based on Models of the Atom
http://baestudent.s-cool.co.uk/animations_interactions.asp http://intro.chem.okstate.edu/workshopfolder/electronconfnew.html
Electron Configuration Rules Aufbau principle- Electrons enter orbitals of lowest energy first. For order of orbitals from lowest to highest learn the ZIG ZAG RULE.
Zig-Zag Rule 7s 7p 6s 6p 6d 6f 5s 5p 5d 5f 4s 4p 4d 4f 3s 3p 3d 2s 2p 1s #e- Atomic # 8 92 118 at 7p filled 18 60 zig 6 32 40 zig 5 32 22 zig 4 18 12 zig 3 8 4 zig 2 2 2 zig 1
Electron Configuration of Helium 4s 3 d 3s 3p 2s 2px 2py 2pz 1ss (filled) He 1s 2
Hund s Rule Pauli exclusion principle In a single energy level, electrons must occupy only one orbital until each orbital has an electron. Since electrons have the same charge, they have strong repulsion forces, pushing them into different suborbitals of the same energy. Orbits fill with electrons of opposite spin. (½+ and ½-) Electrons (e-) have strong, negative repulsion forces, but they are less repulsive if they have opposite spin.
Electron Configuration of Nitrogen 4s 3 d 3s 3p Predict the electron configuration of Ne. 2s 2px 2py 2pz 1ss N 1s 2 2s 2 2p 3
Electron Configuration of Neon 4s 3s 2s 2px 2py 2pz 1ss 3 d 3p Predict the electron configuration of Ar. Filled 2 nd energy level (8 electrons = octet) Ne 1s 2 2s 2 2p 6
Electron Configuration of Argon 4s 3s 2s 2px 2py 2pz 1ss 3 d 3p 3rd energy level (8 electrons = octet) Filled 2 nd energy level (8 electrons = octet) Ar 1s 2 2s 2 2p 6 3s 2 3p 6
Observe this Shockwave Electron Configuration which introduces us to the mystery of some periodic trends. http://intro.chem.okstate.edu/workshopfolder/electronconf new.html 1. What happens to the size of the atom as the energy levels are filled? 2. When do the energy levels change the most?
Photoelectric Effect Click here for photoelectric effect simulation High frequency light frees electrons from reactive metals
Periodic Trends Patterns in the physical and chemical properties of the elements are called trends. Periodic means cycle or repeating pattern.
Periodic Trends in Atomic radius Group trends- Radius increases as electrons (must fill new energy levels) and are added to atom. r Atomic mass, and number increase in this direction also.
What happened to energy levels as p+ and e- increased across a row? How does this affect atomic radius across a row?
Periodic Trends in Atomic radius Period trends- Radius decreases as electrons fill across same energy level. Filled inner levels shield outermost electrons from the nucleus (So in any period, between the nucleus and outer electrons, there is the same number of electrons.) This trend is opposite for atomic mass & number.
Electronegativity Metal vs. Nonmetal Where are the metals vs. the nonmetals? Nonmetals have high electronegativity. Metals have low electronegativity
Electronegativity vs. atomic size Where are the smallest atoms in a period? Big atoms have lower electronegativity
Electronegativity explained e- e- Valence electrons of small atoms that are closer to the nucleus than larger atoms, tend to be held to the nucleus with stronger forces of attraction. Usually the farther they are away, the weaker the forces of attraction. High or strong attraction to valence electrons in a bond = High Electronegativity.
First Ionization Energy This is the amount of energy it takes to remove the first or outer most electron. Look on your periodic table at first ionization potential in V, or on page 362-3 in textbook. How easy is it to remove electron from the Group I & II metals? From the halogens?
First Ionization Energy
How might the first ionization energy compare to the electronegativity across the first period?
They are very similar!
Trends Important to Bonding Ionization energy was used to help determine electronegativity. Electronegativity is a scale in the units of Paulings, developed or calculated to show the degree one element tends to have the bonding electron(s) in a pair of oppositely charged ions. NaCl http://jcrystal.com/steffenweber/java/jpt/jpt.html
For example 0.82 3.12 The electronegativity of nonmetal Cl- is 3.12 And the electronegativity of metal K+ is.82 The difference between the two is 2.30 Pauling units We determine the percent ionic character of the bond to be 74%. By definition this is considered ionic bonding since it is more than a 2.0 difference.
To determine bond type: Calculate the difference in Electronegativity between these Element Pairs Difference > 2.0 = ionic Difference < 2.0 = covalent Li and F 3.0 C and H 0.35 Na and F 3.05 N and H 0.94 K and F 3.16 S and H 0.48 (Formula Units of bonded ions) C and F 1.43 (Molecules of bonded atoms)
Write electron configurations of the following: 1. Ca +2 4. Br -1 2. Al +3 5. S -2 3. K 6. N 1. Ar 1s 2 2s 2 2p 6 3s 2 3p 6 2. Ne 1s 2 2s 2 2p 6 3. 1s 2 2s 2 2p 6 3s 2 3p 6 4s 1 4. Kr [Ar] 4s 2 3d 10 4p 6 5. Ar 1s 2 2s 2 2p 6 3s 2 3p 6 6. 1s 2 2s 2 2p 3
Use electron dot diagrams to determine chemical formulas of the ionic compounds formed when the following elements combine. Example: K and I Ca and S +1-1 +2-2