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Atoms and Moles

Section 1 Substances are Made of Atoms

I. Atomic Theory

-As early as 400 B.C., an atomic theory existed that stated that atoms are the building blocks of all matter
-Democritus was the first scientist who believed in atoms (Greek)
-It wasn't until the 1800s that atomic theory was revised based on scientific observations
 A. Law of Definite Proportion
 -Two samples of a given compound are made of the same elements in exactly the same proportions by mass regardless of the sizes or sources of the samples.
-Every molecule of the same type is made of the same number and types of atoms
-Example: Table Salt (Sodium Chloride)
 -consists of two elements in the following proportions by mass:
 -60.66% Chlorine
-39.34% Sodium
 -Every sample of table salt has these same proportions
 B. Law of Conservation of Mass
 -The mass of the reactants in a reaction equals the mass of the products
-Mass cannot be created or destroyed in ordinary chemical and physical changes
 C. Law of Multiple Proportions
 -If tow or more different compounds are composed of the same two elements, the ratio of the masses of the second element (which combines with a given mass at the first element) is always a ratio of small whole numbers.

 II. Dalton's Atomic Theory

 - Dalton revised the early Greek idea atomic theory in the 1800s into a scientific theory that could be tested by experiments
-Has five important principals
-Believed that elements are composed of only one kind of matter and compounds are made of two or more kinds
-Part of his theory that was incorrect is the fact that like atoms can combine with like atoms (such as O2)
-Did not include the fact that atoms are made up of even smaller particles

Dalton's Theory Contains Five Principles
1. All matter is composed of extremely small particles called atoms, which cannot be subdivide, created, or destroyed
2. Atoms of a given element are identical in their physical and chemical 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 form compounds
5. In chemical reactions, atoms are combined, separated, or rearranged but never created, destroyed, or changed
 Section 2: Structure of Atoms
 I. Subatomic Particles
 A. Electrons- negative charge
B. Nucleus- an atom's central region, which is made up of protons and neutrons
 1. Protons- positive charge. Number of protons is atomic number.
2. Neutrons- no charge.
 II. Atomic Number and Atomic Mass
Elements differ from each other in the number of protons their atoms contain
 -Atomic Number- the number of protons in the nucleus of an atom; the atomic number is the same for all atoms of an element
-Atomic numbers are always whole numbers
-Atomic number also reveals the number of electrons in an atom of an element because for an atom to be neutral, electrons must equal protons
 Mass Number is the Number of Particles in the Nucleus
 -Mass number- the sum of the numbers of protons and neutrons of the nucleus of an atom
Example:
mass number- atomic number= number of neutrons
In this example, the neon atom has 10 neutrons
number of protons and neutrons (mass number)= 20
 - number of protons (atomic number)= 10
 number of neutrons= 10
 -Mass as a number can vary among atoms of a single element
- All atoms of an element have the same number of protons but can have different numbers of neutrons
 Example 2: Determining the Number of Particles in an Atom
How many protons, electrons, and neutrons are present in an atom of copper whose atomic number is 29 and whose mass number is 64?
1. Gather info
 -The atomic number of copper is 29
- The mass number of copper is 64
protons= 29                                        64
electrons= 29                                   -29
                                                             35 neutrons
 - Different elements can have the same mass number
- Knowing just the mass number does not help identify the element
Ex: Some copper atom nuclei have 36 neutrons (therefore mass number = 65) Zinc atoms have 30 protons and 35 neutrons
- Isotopes of an Element Have the Same Atomic Number
- Isotope- an atom that has the same number of protons (atomic number) as other atoms of the same element has a different number of neutrons (atomic mass)
- There are two standard methods of identifying isotopes
- Write the mass number with a hyphen after the name of an element (called hyphen notation; ex. Bromine- 80)
- Shows the composition of a  nucleus as the isotope's nuclear symbol
- Nucleic Notation
         Ex. 126C        C= element symbol
                              12= mass number
                                6= atomic number
Chapter 3
Section 3: Electron Configuration
I. Atomic Models
 -After the atomic theory was widely accepted by scientists, models of atoms were constructed.
-Building a model helps scientists imagine what may be happening at the microscopic level
-Models have limitations
-Models are modified or discarded as new information is found
 A. Rutherford's Model Proposed Electron Orbits
  •  From section 2 J.J. Thomson proposed that the electrons of an atom were embedded in a positively charged ball of matter
  • Named the plum-pudding model because it resembled English plum pudding, a dessert consisting of a ball of cake with pieces of fruit in it
  • In 190, Rutherford performed experiments that disproved Thomson's model.
  • Rutherford envisioned the electrons outside the nucleus orbiting like planets orbiting the sun
  • Because opposite charges attract, the negatively charged electrons should be pulled into the positively charged nucleus
B. Bohr's Model Confines Electrons to Energy Level
  •  Rutherford model was replaced two years later by a model developed by Niels Bohr, a Danish physicist
  • According to the Bohr model, electrons can be only certain distances from the nucleus
  • Each distance from the nucleus quantity of energy that an electron can have
  • The distance in energy between two energy levels is known as a quantum of energy
C. Electrons Act Like Both Particles and Waves
  •  Thomson's experiments demonstarted that electrons act like particles that have mass

II. Electrons and Light
  •  By 1900, scientists knew that light could be thought of as moving waves that have given frequencies, speed, and wavelengths
  • Wavelength- the distance between two consecutive peaks or troughs of a wave
    • units- meters
    • wavelength of light- 105 to less than 10-10 m  
  • Electromagnetic Spectrum- all the frequencies or wavelengths of electromagnetic radiation
  • Einstein proposed that light has the properties of both waves and particles
  • Light can be described as a stream of particle, the energy of which is determined by the light's frequency
A. Light is an Electromagnetic Wave
  •  When passed through a glass prism, sunlight produces the visible spectrum--all the colors of light that a human can see
B. Light Emission
  •  When a high-voltage current is passed through a tube of hydrogen gas, lavender-colored light is seen
  • When this light is only made up of a few colors called LINE-EMISSION SPECTRUM
  • Each element has a line-emission spectrum that is made of a different pattern of colors
C. Light Provides Info About Electrons
  •  Ground state- lowest energy state of a quantized system
  • Excited state- state in which an atom has more energy than it does as its ground state
  • If an electron gains energy, it moves from ground state to excited state
Chapter 3: Section 3

III. Quantum Numbers

  •  Quantum model- present-day model of the atom in which electrons are located in orbitals
- Electrons within an energy level are located in orbitals (regions of high probability for findings particular electron)
  • To define the region in which electrons can be found, scientists have assigned four QUANTUM NUMBERS
  • Quantum Number- a number that specifies the properties of electrons
-principal quantum number (n)
-angular momentum quantum number (l)
-magnetic quantum number (m)
-spin quantum number (+  ½ or -  ½) ( ↑ or ↓)
  •  Principal Quantum Number (n) - indicates the main energy level occupied by the electron
-values are positive integers such as 1, 2, 3,and 4
-As n increases, the electron's distance from the nucleus and the electron's energy increases
  •  Angular Momentum Quantum Number (l)- indicates the shape or type of orbital that corresponds to a particular sublevel.
  • Chemists use a letter code for this Number
l=  0 corresponds to an s orbital
l= 1 to a p orbital
l= 2 to a d orbital
l= 3 to an f orbital 
  •  Magnetic Quantum Number (m)- indicates the numbers and orientations of the orbitals around the nucleus
  • The  value of m takes whole-number values, depending on the value of l
  • The number of orbitals includes on s orbital, 3 p orbitals, 5 d orbitals, and seven f orbitals
  • Spin Quantum Number (symbolized by +  ½ or -  ½ and by  ↑ or ↓) - indicates the orientation of an electron's magnetic field
  • A single orbital can hold a maximum of 2 electrons, which must have opposite spins
A. An Electron Occupies the Lowest Energy Level Available
  •  Pauli Exclusion Principle helps you to write an electron configuration for an atom
  • Aufbau Principle- electrons fill orbitals that have the lowest energy first
B. An Electron Configuration is a Shorthand Notation
  •  The arrangement of the electrons can be shown by the nucleus's electron configuration
  • Sulfur has sixteen electrons:
1s2 2s22p6 3s2 3p4


 Section 4: Counting Atoms

I. Atomic Mass 

 - Atoms are so mall that the gram is not a very convenient unit for expressing their masses
-  Atomic Mass- the mass of an atom expressed in atomic mass units (amu)

II.Introduction to the Mole

- Mole- number of atoms in exactly 12 grams of carbon-12. It is the SI unit for the amount of a substance (
- Molar Mass- the mass in grams of one mole of a substance (g/mol)
- Avogadro's number- 6.022  x1023 The number of atoms or molecules in 1 mol.
 3.50 mol Cu x 63.55 x  g Cu                             
                            1 mol Cu

3.50  mol Cu x 63.55 x 1023 g Cu = 222 g Cu
                           1 mol Cu

II. Intro to the Mole

A. Chemists and Physicists agree on a Standard
  •  In 1960, a standard was set based on an isotope of carbon
  • Defines atomic mass unit (amu) as one twelfth of the mass of one carbon-12 atom
  • One amu= 1.6005402 x  10-27 kg
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