1 / 50

Chemistry Review Atomic Theory

Science and Tech 11. Chemistry Review Atomic Theory. ATOMIC THEORY. An atom is the smallest particle of an element that has the properties of that element. ATOMIC THEORY.

donoma
Download Presentation

Chemistry Review Atomic Theory

An Image/Link below is provided (as is) to download presentation Download Policy: Content on the Website is provided to you AS IS for your information and personal use and may not be sold / licensed / shared on other websites without getting consent from its author. Content is provided to you AS IS for your information and personal use only. Download presentation by click this link. While downloading, if for some reason you are not able to download a presentation, the publisher may have deleted the file from their server. During download, if you can't get a presentation, the file might be deleted by the publisher.

E N D

Presentation Transcript

  1. Science and Tech 11 Chemistry ReviewAtomic Theory

  2. ATOMIC THEORY • An atom is the smallest particle of an element that has the properties of that element.

  3. ATOMIC THEORY • An element is a pure substance that cannot be chemically broken down into simpler substances. Example: Oxygen (O) is an element.

  4. ATOMIC THEORY • A compound is a pure substance that is made up of two or more different elements that have been combined in a specific way. • Example: H2O is a compound made of the elements hydrogen and oxygen.

  5. ATOMIC THEORY • An atom includes smaller particles called protons, neutrons, and electrons:

  6. ATOMIC THEORY • Protons are subatomic particles that have a 1+ (positive) charge. • Neutrons are subatomic particles that do not have an electric charge. • Electrons are subatomic particles that have a 1− (negative) electric charge.

  7. THE NUCLEUS • The nucleus is: • at the centre of an atom • composed of protons and neutrons. • Electrons exist in the area surrounding the nucleus.

  8. THE NUCLEUS The atomic number = the # of protons = the # of electrons • the # of protons in the nucleus determine the element.For example, silver (Ag) has 47 protons

  9. The Periodic table of elements

  10. The Periodic table of elementS • The periodic table organizes all known elements in order by atomic number. • Rows of elements (across) are called periods. • Columns of elements (down) are called chemical families or groups.

  11. ROWS = PERIODS COLUMNS= GROUPS

  12. The Periodic table of elementS • All elements in a family have similar properties and bond with other elements in similar ways. • Group 1 = alkali metals • Group 2 = alkaline earth metals • Group 17 = halogens • Group 18 = noble gases • Metals are on the left side of the table, non-metals are on the right side, and the metalloids form a “staircase” toward the right side.

  13. Ion formation • Atoms gain and lose electrons to form bonds. • When atoms gain or lose electrons, they become electrically charged particles called ions.

  14. Ion formation • Metals lose electrons and become positive ions (aka “cation”). • Some metals are multivalent, which means they lose electrons in different ways • Iron (Fe) loses either 2 electrons (Fe2+) or 3 electrons (Fe3+) as shown in the periodic table.

  15. Ion formation • Non-metals gain electrons and become negative ions (aka ‘anions’).

  16. BOHR DIAGRAMS • A Bohr diagram is a model of the atom that describes the arrangement of an element’s subatomic particles.

  17. BOHR DIAGRAMS • Electron shells are regular patterns or energy levels around the nucleus. • There is a maximum of 2 electrons in the 1st shell, 8 electrons in the 2nd shell, and 8 electrons in the 3rd shell. • Electrons in the outermost shell are called valence electrons.

  18. BOHR DIAGRAMs • Patterns seen in the Periodic Table of Elements: • By PERIOD: • The period number equals the # of shells in the atom. • By GROUPS: • Except for the transition elements (Groups 3–12), the last digit of the group number equals the number of electrons in the valence shell.

  19. 1 2 3 4 5 6 7 Example: Period 2 elements have only 2 shells (or energy levels)

  20. Example: Group 1 = 1 valence electron Group 13 = 3 valence electrons

  21. BOHR DIAGRAMs • A full valence shell is very stable and atoms will lose or gain electrons to accomplish this • Why atoms may bond or react with other atoms • The noble gas elements already have full electron shells and are very stable and thus, rarely react with other elements

  22. BOHR DIAGRAMS • Electrons can exist singly as unpaired electrons, or they can be in pairs, called paired electrons.

  23. Science and Tech 11 Chemistry Review (part 2):Forming Compounds

  24. Background: • Atoms want to achieve a full valence shell and will lose or gain electrons to do this. • very stable and have low energy states (ie: noble gases)

  25. Background: • Cations are: • an atom of metal that loses electrons to other atoms • Anions are: • an atom of non-metal that gains electrons from other atoms

  26. Background: • Atoms can also share electrons

  27. Background: • There are two ways that ions can form compounds: • IONIC BONDING (IONIC COMPOUNDS) • COVALENT BONDING (COVALENT COMPOUNDS)

  28. IONIC COMPOUNDS • Ionic compounds are formed usually between a METAL/CATION and a NON-METAL/ANION • one or more electrons are transferred from the metal to the non-metal

  29. Example #1:Sodium (Na) is a metal that can get a full valence shell by losing 1 electron and chlorine (Cl) is a non-metal that can get a full valence shell by gaining 1 electron. They’re perfect for each other and they form sodium chloride (NaCl)!

  30. Example #2:Calcium fluoride (CaF2) is an ionic compounds formed from 1 atom of the metal calcium and 2 atoms of the non-metal fluorine. The calcium atom transfers 2 electrons in total (1 electron to each of the fluorine atoms).

  31. COVALENT COMPOUNDS • covalent compounds/molecules are formed when two non-metal atoms share electrons (“CO” = together/share and “VALENT” = valence electrons  shared valence electrons)

  32. COVALENT COMPOUNDS • the term ‘molecule’ applies specifically to covalent compounds • some common covalent compounds are water, methane, ammonia, and carbon dioxide

  33. COVALENT COMPOUNDS • Bonding pair: • Each atom in a covalent molecule will share 1 valence electron • Any pairs of valence electrons NOT shared are called lone pairs.

  34. COVALENT COMPOUNDS • Some elements are more stable when paired as two atoms that share electrons. These molecules are called diatomic molecules.

  35. COVALENT COMPOUNDS • Elements that are diatomic are: hydrogen (H2), nitrogen (N2), oxygen (O2), chlorine (Cl2), fluorine (F2), bromine (Br2), and iodine (I2).

  36. LEWIS DIAGRAMS • Lewis diagrams are great for illustrating chemical bonding because they show only the valence electrons of atoms • A lot less work than drawing Bohr diagrams! Ca

  37. To draw a Lewis diagram, follow these steps: • figure out how many valence electrons an atom has • place 1 electron at a time starting at the North position until all electrons have been place • you may end up with paired and unpaired electrons

  38. To draw a Lewis diagram to illustrate ionic bonding: • find the electric charge of the element • the cation will have no valence electrons and the anion will have a full valance shell • add square brackets around each atom • include an ion charge to the top right

  39. Example: Sodium chloride (NaCl)

  40. Example: Barium bromide (BaBr2)

  41. To draw a Lewis diagram to illustrate covalent bonding: • identify bonding pair (shared electrons) • draw a line to represent the bonding pair

  42. Example: Methane (CH4) Example: water (H2O)

  43. EXIT SLIP!!

More Related