Chemistry 1 Chapter 5 – Atomic Structure & The Periodic Table

5.1 Early Models of the Atom Dalton’s Atomic Theory

1. All elements are composed of tiny indivisible particles called atoms. 2. Atoms of the same element are identical. The atoms of any one element are different form those of any other element.

3. Atoms of different elements can physically combine with one another in simple whole-number ratios to form compounds. 4. Chemical reactions occur when atoms are separated, joined, or rearranged. Atoms of one element, however, are never changed into atoms of another element as a result of a chemical reaction.

Atom – the smallest particle of an element that retains the properties of that element. Example: How many atoms are in each compound? CO 2 H 2 OH 2 SO 4 337

5.2 Structure of the Nuclear Atom Electron Cathode Ray Tube (CRT) - negatively charged subatomic particle. They have a –1 charge and are 1/1840 the mass of a hydrogen atom - device used to originate a flow of charged particles

CRT

J.J. Thomson - credited with discovering the electron He used CRT's to measure deflection of charged rays using different gasses, magnets, and different metal electrodes. Found that the ray moved from the anode (-) to the cathode (+) He proposed that the ray was composed of negatively charged particles which he called electrons. He also calculated the charge-to-mass ratio of the electron.

Robert Millikan – His studies led to the determination of the properties of the electron in 1916.

E. Goldstein – discovered positively charged particles in cathode ray tubes. –These particles were later called protons.

Proton – positively charged subatomic particle. Neutron - discovered by James Chadwick Has approximately the same mass as a proton. They are located in the nucleus. It has no charge.

Once subatomic particles had been discovered, Dalton’s model of the atom had to be modified.

Thomson proposed the “Plumb Pudding” Model of the atom.

Rutherford – credited with discovering the nucleus. Gold Foil Experiment

Nucleus – the central core of an atom and is composed of protons and neutrons. It accounts for more than 99% of an atoms mass. The nucleus is also positively charged.

Niels Bohr Niels Bohr applies quantum theory to Rutherford's atomic structure by assuming that electrons travel in stationary orbits defined by their angular momentum. This led to the calculation of possible energy levels for these orbits and the hypothesis that the emission of light occurs when an electron moves into a lower energy orbit.

5.3 Distinguishing Between Atoms Atomic Number – the number of protons in the nucleus of an atom.

Example: What is the atomic number for the following elements? AlC Br I Na Ag How many protons are there in each? The SAME# of protons = Atomic #

Mass Number Isotopes – the number of protons and neutrons in the nucleus of an atom. – atoms that have the same number of protons but different numbers of neutrons.

Example: What is the number of neutrons in each of the following isotopes? 27 Al 12 C 80 Br 127 I 23 Na 108 Ag

Find the missing values for the NEUTRAL atoms in the table below.

ElementAtomic NumberMass Number# protons# electrons# neutrons K

Atomic Mass – a weighted average mass of the atoms in a naturally occurring sample of the element. Atomic Mass Unit (amu) – is defined as one-twelfth the mass of a carbon-12 atom.

Example: A sample of cesium is 75% 133 Cs, 20% 132 Cs, and 5% 134 Cs. Calculate its average atomic mass?.75x133= x x = = amu

Compare this value to that of the periodic table. Is it accurate? amu Calculated Value amu Periodic Table I would say…YES!

The Periodic Table The Periodic Table was first put together by Demitri Mendeleev.

Mendeleev’s Table (~1872) Arranged according to increasing atomic mass Elements with similar properties placed in the same vertical column Missing elements were yet to be discovered and were left blank

Henry Moseley – ( ) British Physicist determined the atomic numbers of the elements He arranged the periodic table in order of atomic number His periodic table is the one used today

Periodic Law – When the elements are arranged in order of increasing atomic number, there is a periodic repetition of their physical and chemical properties.

Periods – horizontal rows of the Periodic Table

Groups – vertical columns of elements on the Periodic Table

Representative Elements Group A elements

Alkali Metals – Group 1-A elements Alkaline Earth Metals – Group 2-A elements Halogens – Group 7-A elements Noble Gases – Group 0 elements

Group B Elements Transition Metals – In the middle of the Periodic Table Inner Transition Metals (rare earth metals) – bottom of the Periodic Table

Classification of Elements Metals – found on the left-side of the Periodic Table Metalloids (or semi-metals) – along the stair- step line Properties are intermediate between metals and nonmetals Nonmetals – found on the right-hand side of the Periodic Table