Atoms smallest particle that an element can be

Atoms smallest particle that an element can be divided into. u Atoms comes from the Greek word atomos which means “invincible. ” u Atoms are made up of even smaller particles called protons, neutrons, & electrons. u Protons & neutrons stick together to form the nucleus, which is the center of the atom. u Electrons are located surrounding the nucleus in the electron cloud. u

Dalton's Atomic Theory (1803) • All matter is composed of extremely small particles called atoms. • Atoms are the building blocks of matter. • All atoms of a given element are identical in size, mass, and other properties. Atoms of different elements differ in size, mass, and other properties. • Atoms cannot be subdivided, created, or destroyed. • Atoms of different elements can combine in simple, whole-number ratios to form compounds. • In chemical reactions, atoms are combined, separated, or rearranged.

Modern Atomic Theory 1. Elements are composed of atoms, which are extremely small. 2. All atoms of a given element have the same chemical properties and contain the same number of protons. 3. Compounds are formed by the chemical combination of two or more different kinds of atoms. (Whole numbers only) 4. Atoms are the units of chemical change.

Atoms are composed of three types of particles Particles Symbol Charge Location Protons p+ Positive charge Found in the nucleus Neutrons N 0 Neutral (no) charge Found in the nucleus Electrons e- Negative charge Found orbiting the nucleus

u u u u John Dalton proposed that all substances are made up of atoms. These atoms are small dense spheres that cannot be created or destroyed. J. J Thompson proposed that atoms were made up of small particles called protons (+) & electrons (-). However J. J did not know where the p+ or e- where located. He thought they we are evenly distribute Earnest Rutherford proposed that an atom had a dense positively charged nucleus surrounded by electrons (-) Niels Bohr 2 years after Rutherford proposed his theory Bohr stated that electrons revolved around the nucleus in circular paths called ORBITS. Bohr also stated that there could only be a certain amount of electrons in each orbit.

Brief history of atomic Theory development Year Milestone Scientist(s) The first coherent atomic theory Democritus 1804 First "modern" atomic theory J. Dalton 1869 First periodic table D. Mendeleev Radioactivity discovered and identified H. Becquerel, M. Curie, E. Rutherford Electron discovered J. J. Thomson Identification of an atomic nucleus E. Rutherford, H. Geiger 1913 Mass of electron determined R. A. Millikan 1913 Positive charge in the nucleus and naturally occuring isotopes discovered J. J. Thomson 1913 Atomic number determined, periodic table reorganized H. Moseley 1919 Proton existance confirmed, neutron proposed E. Rutherford 1931 Neutron Identified J. Chadwick ~400 BC 1896 -9 1897 1909 -11

All matter is made of different kinds of atoms. u Different atoms are called ELEMENTS. u Two or more elements can combine to form COMPOUNDS. u Elements are considered the simplest form of matter. u Elements CAN NOT be broken down. u There are 110 elements in the universe. u 92 are found naturally on earth / in its atmosphere. u 18 are made in laboratories. u

u Each element is made of the SAME type of atom. u Example: the element Oxygen is made out of oxygen atoms. u Example: the element hydrogen is made out of hydrogen atoms.

• Earth’s crust is made up mostly of the elements oxygen & silicon • Earth’s atmosphere is about 78% nitrogen & 21% oxygen

Protons + neutrons # of protons

e r e n t n u m b e r o f t h e t i n y e l e c t r o n s , a n d a c o r r e s p o n d i • Every atom is made up of a nucleus (which contains protons and neutrons). • The nucleus is the center, and it contains most of the mass of the atom. • Outside the nucleus there are electrons, which are extremely small, and which move around the nucleus very quickly. • The electrons are segregated into "shells" which are like invisible spheres that go around the atoms. • Only two electrons can be in the innermost shell. The next track will take eight atoms.

Calculating the # of electrons in each shell u u u In other words, electron shells have a limited capacity for electrons. As you might expect, the farther an electron shell is from the nucleus, the larger it is. You can calculate the total capacity of an electron shell using the formula 2 n 2, where n equals the number of the electron shell. For example, for the 1 st electron shell n = 1 and 2 x 1 x 1 = 2, telling us that the capacity of the 1 st shell is 2 electrons as we have already seen. For the 2 nd shell (n = 2) and 2 x 2 x 2 = 8. For an atom to fill its 2 nd electron shell, it would need 10 electrons: 2 to fill the 1 st shell and 8 to fill the 2 nd. The 3 rd shell has a total capacity of 2 x 3 x 3 = 18 electrons.

Shell Number 2 * (the shell number * the shell number) = the maximum number of electrons in this shell. Maximum Number of Electrons in the Shell 1 2 x (1 x 1) = 2 2 x 1 = 2 2 2 x (2 x 2) = 8 2 x 4 = 8 3 2 x ( 3 x 3) = 18 2 x 9 = 18 4 2 x (4 x 4) = 32 2 x 16 = 32 5 2 x (5 x 5) = 50 2 x 25 = 50 There are five possible shells available to atoms. This diagram shows only three shells, and some atoms, such as the little hydrogen atom, use only the one inner shell. Each of the elements has a different number of the tiny electrons, and a corresponding number of protons. Hydrogen is the smallest, with only one electron. Carbon, nitrogen, and oxygen, are also small, light atoms.

Valence electrons u The electrons found on the outermost energy shell of an atom. u The outer shell electrons that are available for and capable of forming bonds with other atoms. Valence electrons are electrons occupying the outermost energy level.

Electron configuration shorthand Hydrogen H )1 e. Lithium Li )2 e- )1 e. Fluorine F )2 e- )7 e. Sodium Na )2 e- )8 e- )1 e-

This is a diagram of an atom of neon gas. I have colored the nucleus orange and the electrons blue. The inner shell is filled by two electrons. The next shell is filled too -- with eight electrons. Neon doesn't need any more electrons to fill its outermost shell, so it doesn't participate in chemical reactions. It is what we call an inert gas, one that doesn't combine with other elements.

Now we see a very different situation. Here we have a carbon atom near two hydrogen atoms. Note that the carbon atom has two electrons in the inner ring: that is filled. But in the next ring, it has four empty spaces! The hydrogen atoms, too, could each use another electron. Now lets look at some more atoms. Here we see that four hydrogen atoms are sharing their electrons with the carbon atom! They have filled the empty ring spaces in the carbon atom, and the carbon atom is sharing its electrons as well. We now have a molecule of methane gas! This is real chemistry!

The materials that the earth are made of can be thought of as being of two kinds: elements and compounds. Elements, such as oxygen (O), gold (Au), iron (Fe), nitrogen (N), and carbon (C), cannot be broken down into anything else. Compounds, such as carbon dioxide, salt, and water, are made of combinations of the elements. Their molecules can be separated back into the elements that they are made of, or other, but simpler, compounds.

Chemical Bonding u Two types of bonding – Ionic bonding – Covalent bonding u Ionic bonding occurs when 1 atom donates 1/more electrons to another atom – DONATE electrons u Covalent bonding atoms combine to share electrons to fulfill the octet rule. – SHARE electrons

Ionic Bonding In ionic bonding, electrons are completely transferred from one atom to another. In the process of either losing or gaining negatively charged electrons, the reacting atoms form ions. The oppositely charged ions are attracted to each other by electrostatic forces, which are the basis of the ionic bond. For example, during the reaction of sodium with chlorine: sodium (on the left) loses its one valence electron to chlorine (on the right), Resulting in a positively charged sodium ion (left) and a negatively charged chlorine ion (right).

Notice that when sodium loses its one valence electron it gets smaller in size, while chlorine grows larger when it gains an additional valence electron. This is typical of the relative sizes of ions to atoms. Positive ions tend to be smaller than their parent atoms while negative ions tend to be larger than their parent. After the reaction takes place, the charged Na+ and Cl- ions are held together by electrostatic forces, thus forming an ionic bond. Ionic compounds share many features in common: Ionic bonds form between metals and nonmetals. In naming simple ionic compounds, the metal is always first, the nonmetal second (e. g. , sodium chloride). Ionic compounds dissolve easily in water and other polar solvents. In solution, ionic compounds easily conduct electricity. Ionic compounds tend to form crystalline solids with high melting temperatures.

Ionic Bonds Where covalent bonds involve shared electrons, "ionic bonds" are formed when one atom gives up an electron from an outer shell (orbital) and the other atom adds the free electron to its outer most orbital, thereby holding the atoms together in an energetically stable unit. When an atom loses an electron it would have one more positively charged proton (+) then electrons, therefore, the atom would be carrying an overall net charge of (1+). When an atom gains and electron it contains one more electron than protons and therefore would be carrying a net charge of (1 -). Atoms which have gained or lost electrons are called ions. Ions are charged, atoms or molecules. Anions carry a negative charge eg. (Cl-) while cations carry a positive charge (Na+). The strong attraction between positive and negative ions gives rise to and ionic bond. For example the ionic bond between sodium and chlorine (sodium chloride, common table salt).

Table salt Sodium has only one electron in its outer-most orbital, therefore, it can achieve the more energetically stable arrangement of a "full outer" shell by losing the lone outer electron. Chlorine on the other hand has seven electrons in its outermost shell, and it can achieve the energetically stable arrangement of a full outer shell by gaining the electron from sodium. Therefore, the atoms of sodium become positively charged cations and the atoms of chlorine become negatively charged anions and these are held together by ionic bonds as "sodium chloride" (Na. Cl) (table salt).

Covalent Bonding The second major type of atomic bonding occurs when atoms share electrons. As opposed to ionic bonding in which a complete transfer of electrons occurs, covalent bonding occurs when two (or more) elements share electrons. Covalent bonding occurs because the atoms in the compound have a similar tendency for electrons (generally to gain electrons). This most commonly occurs when two nonmetals bond together. Because both of the nonmetals will want to gain electrons, the elements involved will share electrons in an effort to fill their valence shells. A good example of a covalent bond is that which occurs between two hydrogen atoms. Atoms of hydrogen (H) have one valence electron in their first electron shell. Since the capacity of this shell is two electrons, each hydrogen atom will "want" to pick up a second electron. In an effort to pick up a second electron, hydrogen atoms will react with nearby hydrogen (H) atoms to form the compound H 2. Because the hydrogen compound is a combination of equally matched atoms, the atoms will share each other's single electron, forming one covalent bond. In this way, both atoms share the stability of a full valence shell.

The unfilled orbitals of four hydrogen can form four covalent bonds by a sharing of pairs of electrons between carbon and hydrogen.

Atoms that have the same number of protons but different numbers of neutrons are called isotopes. The element hydrogen, for example, has three commonly known isotopes: protium, deuterium and tritium.

CW: FCAT Book Pp. 65 -66 u In your NOTEBOOK u Answer questions 30 -34 u Write out questions/tables u Work on your own QUIETLY!!!!! u Will be discussed and graded