Chapter 6: Atoms, Matter, and Substances Flashcards
Define atom.
The smallest part of an element, composed of a nucleus of protons and neutrons surrounded by a cloud of electrons located in orbitals according to their energies.
Define molecule.
A chemically bonded cluster of atoms; the smallest particle in compounds formed by covalent bonding.
Desribe the atomic model of Democritus.
Democritus proposed that everything was made of tiny, indivisible particles. His idea was that the properties of substances were due to characteristics of the atoms they are made from. i.e. metal atoms were hard and strong, water molecules were wet and slippery, etc.
Desribe the atomic model of John Dalton.
Dalton’s atomic model consists of five principles. The impressive thing is that today the last three points are regarded as correct, and the first two are at least partially correct:
1.
All substances are composed of tiny, indivisible particles called atoms.
2.
All atoms of the same element are identical.
3.
Atoms of different elements have different weights.
4.
Atoms combine in whole number ratios to form compounds.
5.
Atoms are neither created nor destroyed in chemical reactions.
Describe the atomic model of J. J. Thompson.
Thomson developed what we call the “plum pudding” atomic model, which envisions atoms as tiny clouds of massless, positive charge sprinkled with thousands of negatively charged electrons.
Describe the atomic model of Ernest Rustherford.
Rutherford’s model includes three key points:
1.
The positive charge in atoms is concentrated in a tiny region in the center of the atom, which Rutherford called the nucleus.
2.
Atoms are mostly empty space.
3.
The electrons, which contain the atom’s negative charge, are outside the nucleus.
Descirbe the atomic model of Neils Bohr.
Bohr’s “planetary model” depicts electrons orbiting the nucleus of an atom like planets. The model allows a maximum of two electrons in the first energy level, eight electrons in each of the second and third energy levels, and higher numbers in higher levels. Orbits that are farther out from the nucleus are for electrons possessing higher energy.
Describe the quantum model of the atom.
The quantum model includes Bohr’s discovery that all electrons possess certain energies. But in the quantum model, the energy is quantized, meaning that energy comes in units. This means that electrons cannot possess any amount of energy: it must be a specific multiple of the smallest unit of energy that can exist, the quantum.
Describe J. J. Thompson’s Cathode ray tube experiment.
Thomson conducted the cathode ray tube experiment in 1897. In this experiment, he placed electrons from a high-voltage source inside a sealed-glass vacuum tube. He placed the electrodes of another voltage source inside the tube, above and below the cathode ray, and discovered that the beam of electrons deflected toward the positive electrode when this voltage was turned on. The deflection led Thomson to theorize that the beam was composed of negatively charged particles, which he called “corpuscles” (electrons). He confirmed that the ray was negatively charged and was able to determine the charge-to-mass ratio of the individual electrons.
Desribe Robert Millikan’s “Oil Drop Experiment.”
Millikan conducted the “oil drop experiment” in 1911. Inside a metal drum, Millikan placed a pair of horizontal metal plates connected to a high-voltage source, with the upper plate connected to the positive voltage and the lower plate connected to the negative voltage. He used an atomizer spray pump to spray a mist of oil above the positive plate, and the droplets acquired a negative charge due to static electricity. As the oil droplets passed through the hole in the upper plate, Millikan could adjust the voltage and make the negatively charged droplets hover, which allowed him to determine the charge of a single electron.
Describe Rutherford’s Gold Foil Experiment.
In Rutherford’s gold foil experiment, he created a beam of α-particles and aimed the beam at a thin sheet of gold foil surrounded by a ring-shaped screen. Most of the particles went straight through the foil and appeared on the screen opposite the source of the particles; however, some particles deflected with a large angle. Because of Thomson’s theory that the positive charge in the atom was dispersed throughout the atom, Rutherford expected all of the positively charged α-particles to go straight through the foil, so the results of the experiment were astonishing.
How many cm^3 in a Liter?
1 L = 1000 cm^3
How many L in a m^3?
1 m^3 = 1000 L
How many mL in a cm^3?
1 cm^3 = 1 mL
