Unit 2 Study Guide Chemistry

Unit 2 Study Guide 2.1 Emergence of Atoms Law of Conservation of Mass- Mass is never lost or gained during a chemical and physical change. If you have a open experiment, mass can be lost as gas and escape. But if you have a closed experiment, you will keep all the mass. Law of Definite Proportions- Chemical compound contains same elements in same proportions by mass regardless of size per sample Law of Multiple Proportions- If 2 or more compounds have the same elements then the ratio of the masses of the 2nd element combined with the 1st element is the ratio of small whole numbers. Ex. 2.66 oxygen 1 carbon 1.33 g oxygen 1 carbon Ratio is 2:1 Mass Ratio is ratio of one element to another in a compound. 1. All elements composed of small particles called atoms a. Law of conservation of mass 2. Atoms of a given element have the same size, mass, and properties. Atoms of different elements have different properties. a. Law of Definite and Multiple Proportions. i. Same elements have the same mass 3. Atoms cannot be broken down a. Law of Definite and Multiple Proportions 4. Atoms combine to form compounds a. Law of Definite and Multiple Proportions 5. Atoms can be combined, separated, or rearranged in chemical reactions a. All Laws Daltons Atomic Theory 1. All elements are composed of extremely small particles called atoms. a. Now we know there are smaller particles WITHIN atoms called corks ect 2. Atoms of a given element are identical in size, mass, and other properties; atoms of different elements differ in size, mass, and other properties. a. True 3. Atoms cannot be subdivided, created, or destroyed. a. Now we can do that to make bombs 4. Atoms of different elements combine in simple whole-number ratios to form chemical compounds. a. True 5. In chemical reactions, atoms are combined, separated, or rearranged a. True

Thompson’s cathode ray tube Cathode rays were deflected when a negative charge was put near it. This showed that electrons were negatively charged. Plum Pudding Model A sphere with a positively charged pudding. Had negatively charged electrons randomly dispersed throughout the pudding. Rutherford’s Gold Foil Experiment Alpha particles were shot at gold foil. Almost always went through, but very rarely reflected and bounced off. This led to the discovery of the very small nucleus within an atom. This was positively charged because when the two positively particles collided, the beam bounced off. This led to the discovery that there was a dense and small nucleus and the electrons were spread around it. Robert Millikan’s Oil Drop Experiment He had oil drops get pulled by a known gravitational force. He also had a magnetic field pushing the oil drop upward. When the oil drop was levitating in mid air he knew the magnetic force and gravitational pull were equal. He then was able to solve for the charge of the electron.

2.2 Subatomic Particles and Isotopes Rutherford found through his gold foil experiment that there was a dense positively charged nucleus. Thompson found out that there were negatively charged particles that revolved around the nucleus. The placement of the electrons of an element can be used to identify a number of different properties including the number of orbitals, the energy level, and the atomic number. Neutrons- They are place holders so the positively charged protons don’t touch each other and bounce off. Hold the nucleus together. Atomic Number- the number of protons of a given element. Atomic Mass is the mass of the atom. Atomic mass – atomic number is the number of neutrons. Isotope- atoms of same element with different number of neutrons. The atomic mass on the periodic table is affected by the isotopes because it is an AVERAGE of the naturally occurring isotopes. That’s why its normally a decimal. A weighted average means rather than all the data contributing equally to the final average, some data points contribute more than others. Ex.

The percent abundance for chlorine-35 is 75.53%. The percent abundance for chlorine-37 is 24.4%. The mass for Cl-35 is 35.0 amu and for Cl-37 it is 37.0 amu.

2.3 The nature of light, niels bohr, and quantum mechanics Wave–particle duality-the concept that all energy (and thus all matter) exhibits both wave-like and particle-like properties Different Types of Electromagnetic Energy In wave form: • • • • •

x rays ultraviolet light infrared light microwaves Radiowaves

The difference of the wave forms depends of the frequencies of wavelengths. Shorter wavelength and faster frequency correspond to higher energy levels. Visa Versa. In order of where they lie on the spectrum from least the greatest • • • • •

Radio waves microwaves infrared light ultraviolet light x rays

Black Body Radiation- Where solid or liquid is heated and emits light. Like a toaster getting red from being so hot. Planks Constant- ****The energy increased between each level change? The color of the photon depends on how much energy is released, which depends on the frequency (of my tits)

Photoelectric effect- light shone on metal and electrons were produced. Strength of light had no effect but the rate at which they were released had an effect. So good

we have l=h/mv where mv is mass times velocity which is equal to momentum. This allows us to calculate the wavelength of a particle in motion., weird l thing is equal to wavelength and h is plank’s constant white light- combination of all colors colored light- specific light on the spectrum black absorbs light, white reflects the atom’s electrons have their own unique frequency, giving off a different color of light or energy Bohr realized thatwhen energy was added the electron became excited and when the energy was released so was a photon Different orbitals in Bohr model, logic is that there are specific orbitals for each electron Ground state- lowest energy state of atom Excited state- higher than ground state There is a different emission line spectrum for each element because they have unique numbers of electrons and different energies Each element has its own specific spectrum, so an element can be identified by the spectrum, ex. Fingerprint

2.4 Electron configurations Bohr’s model is simplified because now we know the different shapes of each orbital, which are not always in an oval Maximum number of electrons per energy level is 2n2 where n is the energy level Hund’s Rule- electrons would rather be alone if another equal energy orbital is available Aufbau Principle- electrons go to the lowest energy orbital first Pauli Exclusion Principle- every orbital can hold 2 electrons if they have opposite spins Higher atomic number means more potential energy