Ch 3 Atoms and Moles

1. Ch 3 Atoms and Moles

2. Atoms • Atomic theory: Democritus (400 BC) atomos. Indivisible. All substances made of atoms. • Experiments support laws: • Law of definite proportions: • Chemical compounds always contain the same elements in exactly the same proportion by mass • Law of conservation of mass: • Mass can not be created or destroyed in ordinary chemical and physical changes. • Law of multiple proportions: • Two elements combine to form two or more compounds with a given mass of the other in ratios of small whole numbers.

3. Daltons atomic theory (1808) • Using atomic concepts & previous pg. Laws • 1. All matter is composed of extremely small particles called atoms • 2. Atoms of a given element are identical in their chemical and physical properties • 3. Atoms of different elements differ in their physical and chemical properties. • 4. Atoms of different elements combine in simple whole number ratios to from cpds. • 5. In chemical reactions, atoms are combined, separated or rearranged but never created, destroyed or changed.

4. Revisions • #1 revised due to nuclear reaction. #2 revised due to isotopes. • Diatomic molecules H2, O2 etc. • Subatomic particles: proton, neutrons and electrons.

5. Structure of the atom • Experiments to identify subatomic particles • Electrons: • JJ Thomson using cathode rays. Flow of particles away from negative towards positive. • Plum pudding model • Charge of an electron -1.602 x 10-19 C • Mass of an electron 9.1 x 10-31 kg + + + +++ ---- - - + -------------- anode - cathode + <-- flow Plum pudding model

6. Structure of the atom • Rutherford • Gold leaf experiment • Alpha particles directed towards gold foil • Results • Most past straight through, some where deflected back (1/20000) • Revised model: • Small dense positive charged nucleus, volume defined by locations of electrons, atom mostly empty space +

7. Nucleus • 1/10000 the radius of the atom • Protons: positive charge and 2000 time the mass of an electrons • Neutral atom: # protons = # electrons • Chadwick discovered neutron, no charge, mass ~ proton, located in nucleus.

8. Nucleus • Particle symbol charge/notation mass • Proton 1+1 p +1.6x10-19C/ +1 1.673x10-27kg • Neutron 10n 0/0 1.675x10-27kg • Protons in the nucleus repel (like charges) each other • Neutrons are used to balance the repulsive force to stabilize the nucleus,

9. Atomic number and mass number • Atomic number: • # of protons in the nucleus, • Same # for all atoms of the same element • Equal the # of electrons in a neutral atom. • Mass number: • Sum of the number of protons and neutrons of an atom • Maybe different for isotopes of the same element. • Atomic structure/symbol. A = mass number, Z = atomic number X = symbol of element 1H2H3H 11 1 AX Z Protium deuterium tritium

10. Isotopes • Atoms with same # of protons but different # neutrons. Different mass #, same atomic number. • Fill in the blanks: • Element symbol atomic # mass # #p #n • 11 5 • oxygen 16 • 7 8

11. Electron configuration • Atomic models: • Rutherford: positive dense nucleus and electron cloud. • Bohr: positive dense nucleus, electrons in specific dist. and energies. Different energies called quantum of energy. • Louis de Broglie: electrons move as waves • Constructive interference and diffraction patterns. • Electrons location in orbital (clouds) regions in an atom where there is a high probability of finding the electron.

12. Electrons and light • Light: move as waves that have given frequencies, speeds and wavelength. • Speed( c)= 3.0 x 108m/s, constant • Wavelength (): distance between peaks, meters • Frequency (): # waves per second, Hz or 1/s • c= as goes updoes down • Electromagnetic spectrum: range of energy,  High E& low E& short  long  Gamma x-rays ultra violet visible infrared micro TV/radio 1pm, 1020 Hz vibgyor 10km, 100kHz

13. Particle properties of electrons • Photoelectric effect: • Solar calculator • Minimum frequency to move electrons • Light emission: • Added energy to atom, electron moves from grounded state to excited state (only certain dist.) • Electron falls down to lower state and releases energy of specific  • If in the visible light spectrum --> see color. • Grounded state: lowest energy state of a quantized system • Excited state: state at which an atom's electron; has more energy than the grounded state

14. Electron configuration • Describes to location (probable) of each electron in an atom (grounded state) using quantum numbers. • Quantum numbers/model • Specific properties of electrons in an atom • Electrons within an energy level are located in orbital (regions of high probability of find an electron) • Four quantum numbers to describe location and properties of electrons

15. 4 quantum numbers • Principle quantum number (n): • Main energy level • Positive integer • As n increases so does distance from the nucleus and energy of electron. • Angular momentum quantum number (l): • Sublevel of main energy level • Indicated the shape of the sublevel (pg 96, figure 21) • l = n-1 • If n = 1 then l = 0, s orbital and spherical shape • If n = 2 then l = 1 or l = 0, l=1 p orbital peanut shaped • If n = 3 then l = 2, l =1 or l = 0, l = 2 d orbital double peanut shaped • if n = 4 then l = 3, l = 2, l = 1 or l = 0, l = 3 f orbital flower shaped.

16. Quantum numbers continue • Magnetic quantum numbers (m): • Indicate the number and orientation of orbital around the nucleus. • If l = 0 the m = 0, one orbital (s) • If l = 1 then m = -1, 0, +1, three orbital (p) • If l = 2 then m = -2, -1, 0, +1, +2, five orbital (d) • If l = 3 then m = -3, -2, -1, 0, +1, +2, +3, seven orbital (f) • Spin quantum number(ms) • + 1/2 • Indicates the orientation of am electron’s magnetic field • 2 electrons w/in an orbital have opposite spin to balance their magnetic fields

17. Electron configuration (short hand) • Location of each electron with in an atom • Rules: • Pauli exclusion principle: no two electrons in the same atom will have the same 4 quantum numbers, will always have opposite spin if in the same orbital. • Hund’s rule: atom in the grounded state, electrons in a sublevel will fill each orbital before pairing up • Aufbau principle: electrons fill orbital from lowest energy level to highest. • Use periodic table to determine order of fill • Orbital notation: all 4 quantum numbers listed.

18. Counting atoms • Atomic mass: mass of an atom expressed in atomic mass units, amu (u) • Mass of p & n ~ 1 amu • Atomic weight: relative average of all the isotopes of an element (g/mol) • Mole (n): SI unit used to measure the amount of a substance. # of particles is the same as the number of atoms in 12.00g of carbon-12 • Molar mass: mass in grams of 1 mol of a substance = to atomic weight of the element. • Avogadro's number: 6.02 x 1023, number of atoms, ions, molecules or particles in 1 mole • 1 mole = 6.02 x 1023 • 1 mole = molar mass.

19. Problems • Calculate the mass of 5.50 mol of sodium. • How many moles of helium are contained in 0.255g He? • How many moles are represented by 8.0 x 1013 atoms of Ca? • Calculate the mass of 4.89 x 1022 atoms of Zn. • How many atoms are represented in 75.5g Ba?

20. Molar mass of cpds • Sum of all the atomic mass within a cpd • Example NaCl, BaSO4 • Convert 3.91g NaCl to number of particle NaCl • Convert 8.83 x 1024molecules CO2 to mass CO2